Chemistry
Exam Review
Chemistry
• States of Matter
• Chemical vs Physical Reactions
• Scientific Measurement
• Density
• Atomic Structure
• Ionic and Covalent Bonding
• Acids and Bases and pH
• Lab Safety
• Balancing Chemical Equations
States of Matter
Gas
Evaporation
Sublimation
Solid
Condensation
Liquid
Melting
Freezing
Chemical vs Physical Reactions
• Chemical Reaction requires at least one of the following:
•
•
•
•
•
Change in Temperature
Change in Color
Create a New Substance
Gas is formed (bubbles!!)
Change in Smell
• Physical Reaction changes the shape or size but not the chemical
makeup of the substance.
How Scientists Measure Matter
• Scientists are often called on to make
measurements of matter, which may include
such things as mass (weight), volume, and
temperature. A worldwide measurement system
has been adopted to ensure that scientists can
speak the same language.
The SI system of scientific measurement
• The SI system (from the French Systeme International) is a worldwide
measurement system based on the older metric system that most of
us learned in school. There are minor differences between the SI and
metric systems, but they’re basically interchangeable.
• SI is a decimal system with basic units for things like mass, length, and
volume, and prefixes that modify the basic units. For example, the
prefix kilo- (k) means 1,000. So a kilogram (kg) is 1,000 grams and a
kilometer (km) is 1,000 meters.
• Two other very useful SI prefixes are centi- (c) and milli- (m), which
mean 0.01 and 0.001, respectively. So a milligram (mg) is 0.001 grams
— or you can say that there are 1,000 milligrams in a gram.
SI/English conversions
• Many years ago, there was a movement in the U.S. to convert to the metric system. But
Americans are still buying their potatoes by the pound and their gasoline by the gallon.
Most professional chemists use both the U.S. and SI systems without any trouble. It’s
necessary to make conversions when using two systems.
• The basic unit of length in the SI system is the meter (m). A meter is a little longer than a
yard; there are 1.094 yards in a meter, to be exact. But that’s not a really useful
conversion. The most useful SI/English conversion for length is:
• 2.54 centimeters = 1 inch
• The basic unit of mass in the SI system for chemists is the gram (g). And the most useful
conversion for mass is:
• 454 grams = 1 pound
• The basic unit for volume in the SI system is the liter (L). The most useful conversion is:
0.946 liter = 1 quart
Two basic rules are associated with the unit
conversion method:
• Rule 1: Always write the unit and the number associated with the
unit. (Rarely in chemistry will you have a number without a unit. The
mathematical pi is an exception.)
• Rule 2: Carry out mathematical operations with the units, canceling
them until you end up with the unit you want in the final answer. In
every step, you must have a correct mathematical statement.
Scientific Measurement
•Remember:
•Metric System of Measurement
•As exact as possible with no rounding
off or heaping measurements
Density
density=mass/volume
Principle 1: If you pack more mass into a the
same volume, it's more dense.
Principle 2: If you pack the same mass into a
smaller volume, it's more dense.
Principle 3: Just because something has more
mass doesn't mean it's more dense.
Atomic Structure
What are the 3 major parts of an
atom?
• Proton
• Neutron
• Electron
Diagram of the atom.
Diagram showing the charges of each part of
the atom.
Bohr Model
Proton
• Protons are positively charged particles found in the
atomic nucleus. Protons were discovered by Ernest
Rutherford..
• Experiments done in the late 1960's and early 1970's
showed that protons are made from other particles called
quarks. Protons are made from two 'up' quarks and one
'down' quark.
Neutron
• Neutrons are uncharged particles found in the atomic
nucleus. Neutrons were discovered by James Chadwick
in 1932.
• Experiments done in the late 1960's and early 1970's
showed that neutrons are made from other particles
called quarks. Neutrons are made from one 'up' quark and
two 'down' quarks.
Electron
Electrons are negatively charged particles that surround the
atom's nucleus. Electrons were discovered by J. J.
Thomson in 1897.
Electrons determine properties of the atom. Chemical
reactions involve sharing or exchanging electrons.
Nucleus
The nucleus is the central part of an atom. It is composed of
protons and neutrons.
The nucleus contains most of an atom's mass.
It was discovered by Ernest Rutherford in 1911.
Quark
• Believed to be one of the basic building blocks of matter. Quarks
were first discovered in experiments done in the late 1960's and
early 1970's.
• Three families of quarks are known to exist. Each family contains
two quarks. The first family consists of Up and Down quarks, the
quarks that join together to form protons and neutrons.
• The second family consists of Strange and Charm quarks and
only exist at high energies.
• The third family consists of Top and Bottom quarks and only exist
at very high energies.
Isotope
Atoms that have the same number
of protons but different numbers
of neutrons
Isotope
Example:
What is the Electron Cloud Model?
Model of the atom pictures the electrons moving
around the nucleus in a region called an electron
cloud.
The electron cloud is a cloud of varying density
surrounding the nucleus. The varying density
shows where an electron is more or less likely to
be. Atoms with electrons in higher energy levels
have additional electron clouds of different
shapes that also show where those electrons
are likely to be.
Electron Cloud Model
Diagram 1:
Electron Cloud Model
Diagram 2:
Bonding
Ionic
Covalent
(Metallic)
How do atoms bond(join) together
to form the millions of different
compounds that make up the
world?
It all comes down to the
electrons!
There are many ways electrons can interact
with one another:
• Electrons can be transferred between
atoms
Ionic Bonding
• Electrons can be shared between atoms
Covalent Bonding
• Electrons can move freely between atoms
Metallic Bonding
Ionic Bonding
Ionic Bonding
• Metals react with non-metals
• Ions form when metalatoms donate
electrons to a non-metal atom
• Metals form +vecharged ions as they
have lostelectrons
• Non-metals form -vecharged ions as they
have gained electrons
• The positive and negative ions share a
strong electrostatic force of attraction –
IONIC BOND
• Ions bond to form IONIC
SUBSTANCES
Consider reactions between the following
metals and non-metals:
sodium + chlorine  sodium chloride
magnesium + oxygen  magnesium oxide
calcium + chlorine  calcium chloride
1 Sodium atom donates 1electron
1 Chlorine atom accepts 1 electron
This means…
For every 1 Sodium atom –
1 Chlorine atom is needed
Na1+
Cl1-
1 Magnesium atom donates 2electrons
1 Oxygen atom accepts 2 electrons
This means…
For every 1 Magnesium atom –
1 Oxygen atom is needed
Mg2+
O2-
1 Calcium atom donates 2electrons
1 Chlorine atom accepts 1 electron
This means…
For every 1 Calcium atom –
2 Chlorine atoms are needed
Cl1-
Ca2+
Cl1-
Have you noticed how
electrons are arranged
within the shells?
Individually??
In pairs??
In groups??
Ionic Lattice
• Not just a pairof ions
• Many atoms (ions) bond to form an
IONIC
LATTICE
• The number of positive and negative ions vary, depending
on how many electrons are transferred between the metal
and non-metal.
1 positive ion : 1 negative ion
Sodium loses 1 electron
Chlorine gains 1 electron
1 positive ion : 2 negative ions
Calcium loses 2 electrons
Chlorine gains 1 electron
Properties of Ionic Substances
The very strong ionic bonds within the lattice means…
• HIGH melting point (usually over 250oC)
• HARD– when apply force the force is spread throughout
lattice
• BRITTLE– large force can cause ions to move, therefore
repel one another  lattice breaks
Properties of Ionic Substances
The very strong ionic bonds within the lattice means…
• DO NOT conduct ELECTRICITY when solid – ions are
not free to move
• WILL conduct ELECTRICITY when in aqueous
solution– ions are now free to move
Examples of Ionic Substances
Sodium Chloride
Magnesium Oxide
Calcium Chloride
Can you think of some others?...
Covalent
Bonding
Covalent Bonding
• Most compounds (substances) in the
world are formed through covalent
bonding
• Non-metals react with non-metals
• Atoms shareapairofelectrons– this is
called a COVALENT BOND
• Two types of covalent bonds:
Covalent Molecular ANDCovalent
Network
Covalent Molecular
• Electron clouds of
atoms overlapand
interact to share
electrons
• Strong ‘electrostatic’
attraction between the
positive nucleus and
negative electrons –
keeps atoms together
• Form molecules
Cl
Cl
Chlorine has 7
electrons in its
valence shell.
It needs 1 electron
for this shell to be
full.
How will 2
chlorine atoms
react and bond?
Cl Cl
Each atom of Chlorine
achieves a full valence shell
by sharing the 2 electrons
in the middle.
How many bonds?
• Atoms may form multiple covalent bonds - share
not just one pair of electrons, but two or more pairs
• Atoms of different elements will form either one,
two, three or four covalent bonds with other atoms
• Number of covalent bonds is equal to eight minus
the group number:
Group
Exampl
e
4
5
6
7
Carbon Nitrogen Oxygen Fluorine
# of
• Hydrogen
forms
8 – 4 one
= 8covalent
– 5 = 8bond,
– 6 = and
8 – the
7=
Bonds
noble
gases in 4Group VIII
3 do not? form covalent
?
bonds at all.
•Oxygen has 6 electrons in its valence shell, therefore needs
2 electrons for this shell to be full.
•2 Oxygen atoms share 2 electrons (2 bonds)–
this fills both of their valence shells.
•Nitrogen has 5 electrons in its valence shell, therefore
needs 3 electronsfor this shell to be full.
•2 Nitrogen atoms share 3 electrons (3 bonds)–
this fills both of their valence shells.
Arrangement
• Do not form a lattice
• Remain as individual molecules
• Interact weakly with other molecules
• Covalent molecules can be small and
simple
• Covalent molecules can be large and
complex
Properties of Covalent Molecular
Substances
• GASES and LIQUIDS at room temperature –
OR easily melted solids
• LOWmelting points – forces between atoms
are strong, BUT forces between molecules
are weak
• DO NOT conduct electricity – no charged
particles (ions)
Covalent Network
• Do not exist as individual molecules
• Form giant networks of covalently bonded atoms
• Carbon (C) and Silicon (Si)
• Diamonds – carbon atoms form 3D
network
• Silica (found in sand) – silicon and oxygen
atoms form a 3D network
Properties of Covalent Network Substances
• HARD and BRITTLE
• HIGH melting points – strong covalent
bonds between molecules in network
** Diamonds melt at just over 4000oC!
• DO NOTconduct ELECTRICITY – no
charged particles (ions)
• INSOLUBLE in water – the bonds will not
break if you add substance to water
• Non-reactors
Examples of Covalent Substances
Carbon Dioxide
Methane
Diamonds
Can you think of some others?...
Acids vs Bases
• For thousands of years people have known that
vinegar, lemon juice, and many other foods taste
sour. However, it was not until a few hundred
years ago that it was discovered why these
things taste sour – because they are all acids.
The term acid, in fact, comes from the Latin term
acere, which means "sour". While there are
many slightly different definitions of acids and
bases, in this lesson we will introduce the
fundamentals of acid/base chemistry.
• In the seventeenth century, the Irish writer and amateur chemist Robert
Boyle first labeled substances as either acids or bases (he called bases
alkalies), according to the following characteristics:
• Acids taste sour, are corrosive to metals, change litmus (a dye extracted
from lichens) red, and become less acidic when mixed with bases.
• Bases feel slippery, change litmus blue, and become less basic when mixed
with acids.
• While Boyle and others tried to explain why acids and bases behave the
way they do, the first reasonable definition of acids and bases would not
be proposed until 200 years later.
• In the late 1800s, the Swedish scientist Svante Arrhenius proposed
that water can dissolve many compounds by separating them into
their individual ions. Arrhenius suggested that acids are compounds
that contain hydrogen and can dissolve in water to release hydrogen
ions into solution. For example, hydrochloric acid (HCl) dissolves in
water as follows:
• HClH2O→H+(aq) + Cl-(aq)
• Arrhenius defined bases as substances that dissolve in water to
release hydroxide ions (OH-) into solution. For example, a typical base
according to the Arrhenius definition is sodium hydroxide (NaOH):
• NaOH H2O
•→
• Na+(aq) + OH-(aq)
• The Arrhenius definition of acids and bases explains a number of
things. Arrhenius's theory explains why all acids have similar
properties to each other (and, conversely, why all bases are similar):
because all acids release H+ into solution (and all bases release OH-).
The Arrhenius definition also explains Boyle's observation that acids
and bases counteract each other. This idea, that a base can make an
acid weaker, and vice versa, is called neutralization.
Neutralization
• As you can see from the equations, acids release H+ into solution and
bases release OH-. If we were to mix an acid and base together, the
H+ ion would combine with the OH- ion to make the molecule H2O,
or plain water:
• H+(aq) + OH-(aq) → H2O
The neutralization reaction of an acid with a base
will always produce water and a salt, as shown below:
The neutralization reaction of an acid
with a base will always produce water
and a salt, as shown below:
Acid
Base
Water
Salt
HCl
+
NaOH
→
H2O
+
NaCl
HBr
+
KOH
→
H2O
+
KBr
• Though Arrhenius helped explain the fundamentals of acid/base chemistry,
unfortunately his theories have limits. For example, the Arrhenius
definition does not explain why some substances, such as common baking
soda (NaHCO3), can act like a base even though they do not contain
hydroxide ions.
• In 1923, the Danish scientist Johannes Brønsted and the Englishman
Thomas Lowry published independent yet similar papers that refined
Arrhenius' theory. In Brønsted's words, "... acids and bases are substances
that are capable of splitting off or taking up hydrogen ions,
respectively." The Brønsted-Lowry definition broadened the Arrhenius
concept of acids and bases.
• The Brønsted-Lowry definition of acids is very similar to the Arrhenius
definition: Any substance that can donate a hydrogen ion is an acid. (Under
the Brønsted definition, acids are often referred to as proton donors
because an H+ ion, hydrogen minus its electron, is simply a proton).
• The Brønsted definition of bases is, however, quite
different from the Arrhenius definition. The Brønsted
base is defined as any substance that can accept a
hydrogen ion. In essence, a base is the opposite of an
acid. NaOH and KOH, as we saw above, would still be
considered bases because they can accept an H+ from
an acid to form water. However, the Brønsted-Lowry
definition also explains why substances that do not
contain OH- can act like bases. Baking soda (NaHCO3),
for example, acts like a base by accepting a hydrogen
ion from an acid as illustrated below:
• Acid Base Salt
• HCl + NaHCO3 → H2CO3 + NaCl
• In this example, the carbonic acid formed (H2CO3) undergoes rapid
decomposition to water and gaseous carbon dioxide, and so the
solution bubbles as CO2 gas is released.
pH
• Under the Brønsted-Lowry definition, both acids and bases
are related to the concentration of hydrogen ions present.
Acids increase the concentration of hydrogen ions, while
bases decrease the concentration of hydrogen ions (by
accepting them). The acidity or basicity of something,
therefore, can be measured by its hydrogen ion
concentration.
In 1909, the Danish biochemist Sören Sörensen
invented the pH scale for measuring acidity. The
pH scale is described by the formula:
• pH = -log [H+]
• Note: Concentration is commonly abbreviated by using square
brackets, thus [H+] = hydrogen ion concentration. When measuring
pH, [H+] is in units of moles of H+ per liter of solution.
• For example, a solution with [H+] = 1 x 10-7 moles/liter has a pH equal
to 7 (a simpler way to think about pH is that it equals the exponent on
the H+ concentration, ignoring the minus sign). The pH scale ranges
from 0 to 14. Substances with a pH between 0 and less than 7 are
acids (pH and [H+] are inversely related - lower pH means higher [H+]).
Substances with a pH greater than 7 and up to 14 are bases (higher
pH means lower [H+]). Right in the middle, at pH = 7, are neutral
substances, for example, pure water. The relationship between [H+]
and pH is shown in the table below alongside some common
examples of acids and bases in everyday life.
Proper pH is necessary for
efficient digestion; esophagus
pH is 6.8, stomach pH is two,
small intestine pH is eight and
large intestine pH around seven.
What is pH and digestion pH
• pH stands for Potential Hydrogen, degree of concentration of H ions
in the substance or a solution.
• pH value of 0 (strongly acidic) to less than 7 (mild acidic) is acid,
molecules that give off H (hydrogen) maintain an acidic pH
• pH of 7 means neutral
• pH of greater than 7 (mild alkali) to 14 (strong alkali) means base,
molecules that attract H (hydrogen) maintain a basic (alkaline) pH
• pH controls the speed of our body's biochemical reactions.
• Acid pH is hot & fast and alkaline pH is cool & slow.
• What we eat and drink will affect where our body's pH level falls, and
our body's pH will control the activity of every metabolic function
happening in our body.
• pH is behind the body's electrical system and intracellular activity as
well as the way our bodies utilize enzymes, minerals, and vitamins.
The pH varies in the digestive process from stage to stage:
• In the mouth, the pH is in neutral (or close to neutral),
• In the stomach, the pH is acidic at around two.
• In the small intestine, the pH is basic at around 8
• Finally, it reaches seven as it reaches the end (anus).
Lab Safety
• Fire
• Chemical Burns
• Weights
• Safe Handling of Equipment
• Glass
• Metal
• Labels
Balancing Chemical
Equations
Steps of Balancing a Chemical Equation
1.
Identify each element found in the equation. The number of atoms of each type of atom must
be the same on each side of the equation once it has been balanced.
2.
What is the net charge on each side of the equation? The net charge must be the same on
each side of the equation once it has been balanced.
3.
If possible, start with an element found in one compound on each side of the equation. Change
the coefficients (the numbers in front of the compound or molecule) so that the number of
atoms of the element is the same on each side of the equation. Remember! To balance an
equation, you change the coefficients, not the subscripts in the formulas.
4.
Once you have balanced one element, do the same thing with another element. Proceed until
all elements have been balanced. It's easiest to leave elements found in pure form for last.
5.
Check your work to make certain the charge on both sides of the equation is also balanced.
Example of Balancing a Chemical Equation
• ? CH4 + ? O2 → ? CO2 + ? H2OIdentify the elements in the equation: C, H, O
Identify the net charge: no net charge, which makes this one easy!
• H is found in CH4 and H2O, so it's a good starting element.
• You have 4 H in CH4 yet only 2 H in H2O, so you need to double the coeffient of
H2O to balance H.1 CH4 + ? O2 → ? CO2 + 2 H2O
• Looking at carbon, you can see that CH4 and CO2 must have the same
coefficient.1 CH4 + ? O2 → 1 CO2 + 2 H2O
• Finally, determine the O coefficient. You can see you need to double the O2
coefficient in order to get 4 O seen on the product side of the reaction.1 CH4 + 2
O2 → 1 CO2 + 2 H2O
• Check your work. It's standard to drop a coefficient of 1, so the final balanced
equation would be written:CH4 + 2 O2 → CO2 + 2 H2O
Balance the Chemical Equations
• Being able to balance equations is important
because it allows you to predict how a chemical
reaction will occur, including what products will
be created, how much will be produced and how
far a reaction will proceed if you know the
amount of reactants.
What is the balanced equation for:
__ TiCl4 + __ H2O → __ TiO2 + __ HCl
TiCl4 + 2 H2O → TiO2 + 2 HCl
TiCl4 + 2 H2O → TiO2 + 4 HCl
2 TiCl4 + H2O → 2 TiO2 + HCl
TiCl4 + 4 H2O → TiO2 + 4 HCl
What is the balanced equation for:
__ Na3PO4 + __ HCl → __ NaCl + __ H3PO4
Na3PO4 + HCl → NaCl + H3PO4
Na3PO4 +3 HCl → 3 NaCl + H3PO4
3 Na3PO4 + HCl → 3 NaCl + H3PO4
Na3PO4 + 3 HCl → NaCl + H3PO4