Science 10 Review
* Please make notes on your review pkg, or on a separate
sheets of paper to augment the points provided *
CHEMISTRY
REVIEW
What is Chemistry? The branch of science
that deals with changes in matter.
Matter and Types of Matter
Each one of these groups contains elements with similar chemical
and physical properties.
family / group
vertical column with similar
chemical properties
period
horizontal row of elements whose
properties gradually change from
metallic to nonmetallic as you move
from left to right
Metals


Metals makeup more than 75% of the elements
on the periodic table.
Metals are characterized by the following physical
properties.
1. They are shiny ( have a high luster).
2. They are usually solids at room temperature.
3. They are malleable ( can be hammered, pounded, or
pressed into different shapes without breaking).
4. They are ductile (can be drawn into thin sheets or
wires without breaking).
5. They are good conductors of heat and electricity

Metals can be both reactive or inert
◦ Reactive: such as sodium, which will combust into a
flurry of flame when it come on contact with air.
◦ Inert: extremely un-reative, like platinum or gold
Nonmetals






There are 17 nonmetals in
the periodic table,
In general, they can be
grouped together b/c they
DO NOT resemble metal
more than having a
relationship to each other.
They can be a variety of
states, (although they are
usually gases at room
temperature).
They do not have a luster.
They are poor conductors
of heat and electricity.
They generally exist as
molecules.
Metalloids



The metalloids are B, Si,
Ge, As, Sb, Te, Po and At.
The members of this
group are the least
uniform in character.
The metalloids are NOT
as good at conducting as
the metals, but they are
better conductors than
the nonmetals.
◦ Many of them are
known as semiconductors.
Ion Charges
1+
2+
3+
Multiple
charges…check the
box.
3-
2-
1-
Info on the periodic table


Every element on the periodic table is in it’s own box.
Each box has the same information in it.
Atomic Theory
See Atomic Theory Review in your Notes

Every element is made of up of three
subatomic components.
◦ Protons
◦ Neutron
◦ Electrons

Protons and neutrons are in the nucleus in
the middle of the element.

Electrons orbit the outside .These electrons
are drawn to the nucleus because of their
opposite charges
Atomic mass vs. Mass number


Atomic mass is the mass of the whole element.
You add together ALL of its components:
◦ protons + neutrons + electrons = atomic mass
◦ This gives you a number with digits in to the 10,000th
place!
ISOTOPES
Def: Atoms with
differing weights but
are the same type of
elements.
 ex carbon-12 and
carbon-14).



Remember: atoms always
have the same number of
protons. Number of
Protons NEVER
changes!!!
Therefore isotopes have
differing numbers of
neutrons.
ISOTOPES
To make it simpler we use
Mass Number.
◦ protons + neutrons = mass number
◦ Therefore…


mass number – protons = neutrons
mass number – protons = neutrons
Ex. a) vandium-51
◦ mass number – protons = neutrons
◦ 51 – 23 = 28
◦ There are 28 neutrons & 23 protons in vandium.
◦ Try: nickel-58, bromine-79, argon-40, uranium-238
◦
nickel 58:
 mass number – protons = neutrons
 58 – 28 = 30
 There are 30 neutrons & 28 protons in nickel.
◦ bromine-79: 35 protons 44 & neutrons
◦ argon-40: 18 protons & 22 neutrons
◦ uranium-238: 92 protons & 146 neutrons
Ions and the Octet rule
 The octet rule says that atoms tend to gain, lose
or share electrons so as to have eight electrons in
their outer electron shell.
Ex: the halogens--each chlorine is missing only one electron
in their valence orbital, so they each share a valance electron
from each other so the molecules to stabilize.
This is why all the halogens are diatomic!
 Other elements do the same thing.
 In the case of water
Oxygen needs two electrons (to move from 6 to 8
valance electrons)
Each hydrogen has one too many--
 Each
hydrogen gives an electron to the
oxygen; oxygen now has a full valance
orbital
The oxygen shares the electrons with
hydrogen too, so it has a full orbital too.
Formation of ions
Ultimately, elements are lazy! Elements will do
whatever is the easiest way to get a full valance
orbital.
 There are two types of ions:

◦ Cations: give up (lose) electrons. Since they now
have more protons than electrons, they have a
positive charge.
Formation of ions

Sodium has 11 protons and 11 electron; only one electron in its
valance orbital,
(it is much easier for it to lose one than find seven....)

This leaves it with 11 protons and only 10
electrons…thus it now has a charge of 1+ b/c there is
one more proton than electron.
Once the electron is
gone, the orbital
“disappears”

Anions: pick up (gain) electrons. Since they now
have more electrons than protons, they have a
negative charge.
e-
-

Chlorine has 17 protons and 17 electrons; it has seven
electron in its valance orbital
(it is much easier for it to gain one than find seven)

This leaves it with 17 protons and 18 electrons…thus it
now has a charge of 1-, because there is one more
electrons than proton.
Only missing 1
e-
Calculating # of Electrons in Ions

Note: The number of protons in an
atom/ion NEVER change.

# of electron in an ion = # of protons – ion charge
Ex. non-metal anion (O2-):
Atomic # = 8 = 8 protons
# of electron = # of protons – ion charge
# of electron = 8 – (-2)
# of electron in an ion of oxygen is = 10
Ex. metal cation (Ca2+):
Atomic # = 20 = 20 protons
# of electron = # of protons – ion charge
# of electron = 20 – (+2)
# of electron in an ion of calcium is = 18
Maintain the – sign
of the ion charge
Maintain the + sign
of the ion charge
Lets try some!
What would the following elements be as ions?

Magnesium, nitrogen, selenium, iodine, potassium
& oxygen

Magnesium

Has two in it valance
orbital…easier to lose 2 than
gain 6. Become Mg2+

Nitrogen

Has five in it valance
orbital…easier to gain 3 than
lose 5. Become N3-

Selenium

Iodine

Potassium

Oxygen
Bohr Diagrams
Bohr Diagrams are
used to
diagrammatically
represent elements and
ions.
 They show the number
of protons, neutrons
and electrons.


The number of electron
orbitals is equal to the
number of the row the
element is in.
3rd row, 3 circles (orbitals)

The number of electrons
that can fill each orbital is
equal to the number of
elements in each row.
2, 8, 8, 18, 18, etc
1
2
3
4 5 6
7 82

The number of valance
electrons is equal to the
number of the column
that the element is in.
3rd column, 3 valance electrons
1 2
3
To draw a Bohr Diagram
Place a circle in the
center to represent the
nucleus.
 Write the symbol to
represent the element
in the circle.
 Write the number of
protons & neutrons in
the circle to.

Al
13 p+
14 n
Draw in the correct
number of orbital


Recall the number of rings
= column the element is in
The electrons are
placed on the rings

1
2
3 ◦ Recall The number of
electrons that can fill each
orbital is equal to the
number of elements in
each row
◦ AND The number of
valance electrons is equal to
the number of the column
that the element is in.
Al
13 p+
14 n
Types of compound’s

Ionic:
◦ cmpds that have one metal and one non-metal
ion (one positive and one negative)
 Ex. Na+ & Cl- make NaCl(s)

Molecular:
◦ cmpds that have two non-metal ions
 Ex. C & O make CO2 (g)
Ionic Cmpds

Are a result of ionic bonds
◦ Naming: the metal is always first, the non-metal
second. The non-metals name is changed to have an
“ide” ending
 (ie. sodium chloride)

Ionic bonds form between metals and non-metals,
◦ This means that there is a tight arrangement of
particles in rigid pattern, which is hard to break down,
giving them a high melting/ boiling temp.
◦ Conduct electricity (electrolytes)
◦ Solid at room temp
Ionic Bonds


In ionic bonding, valence
electrons are completely
transferred from one atom to
another.
The result? Ions!
◦ Electrically charged atoms.
◦ Cations are positively charged
(Mg 2+, H+, Na+)
◦ Anions are negatively charged
(O2-, Cl-)
The oppositely charged ions are
attracted to each other by
electrostatic forces.

Let's take a look!
How to write Ionic Cmpds:
Step 1
Write each ion with its charge
Step 2
Figure out how many of each you need to
make the charges balance.
Ca2+ and Br –
1 Ca and 2 Br
Ca2+, Br – ,Br –
Step 3
Write the formula using subscripts to show how
many of each atom you need AND the state
Ca2+ and Br –
1 Ca and 2 Br
Ca2+, Br – ,Br –
CaBr2 (s)
* Ionic compounds
are always solid (s)
Multiple charges…
Polyatomic Ions

These are the ions found middle area of
the periodic table
◦ Elements that are already grouped together
and poses a charge.
Molecular Cmpds

Unlike ionic compounds; a positively charged metal ion
and a negatively charges non-metal ion, molecular
compounds are a combination of two non-metals.
◦ Since both are negatively charged, we can not
balance the formula to equal zero.
Covalent bonds



Covalent bonds are formed as a result of the
sharing of one or more pairs of bonding electrons;
(this is what hold molecular compounds together)
Each atom donates half of the electrons to be
shared.
Where the clouds overlap they are thicker, and
their electric charge is stronger.
 See!
Covalent bonds

Molecular Compounds
◦ Do not have a tight crystal structure so the melting/
boiling point is lower.
◦ Do not conduct electricity (non-electrolyte)
◦ Can be any state at room temp.
Molecular cmpds

Naming: The first element in the compound uses the element
name. The second element in the compound has the suffix ide
added to it- just like ionic compounds HOWEVER…
◦
When there is more than one of the atom in the formula, a
prefix is used to specify how many of that element there is.
 mono- 1
hexa- 6
 di- 2
hepta-7
 tri- 3
octa-8
 tetra- 4
nona-9
 penta- 5
deca-10
◦
Here comes the exception…when the first element has
only one, no prefix is used. If the second element has only
one, the prefix mono is attached.

Acids:
◦ Ionic cmpds, where the
metal is always
hydrogen.
◦ In solution will have a
pH lower than 7
◦ React predictably with
indicators like litmus
paper

Bases:
◦ Ionic cmpds, where the
non-metal is always
hydroxide (OH-).
◦ In solution will have a
pH higher than 7
◦ React predictably with
indicators like litmus
paper
Naming Acids

If the compound name ends in “ide” the name of the acid becomes
hydro--ic acid.
◦ For example HCl, hydrogen chloride, become
hydrochloric acid

If the compound name ends in “ate” the name of the acid becomes
--ic acid.
◦ For example H2SO4, hydrogen sulfate, become
sulfuric acid

If the compound name ends in “ite” the name of the acid becomes
--ous acid.
◦ For example HClO3, hydrogen perchlorite, become
perchlorous acid

In general the H+ ion always goes first…unless the
acid has an “organic group”; (a molecule containing
COOH), then the H+ goes last.

Acids are always written with an (aq) subscript.
Lets try some!

RULE 1: hydrogen ---ide
◦ HF(aq)
 Hydrogen flouride
◦ H2P(aq)
 Hydrogen phosphide
◦ HI (aq)
 Hydrogen iodide
becomes
hydro ---ic acid
Hydroflouric acid
Hydrophosphoric acid
Hydroiodic acid
Lets try some!

RULE 2: hydrogen ---ate
◦ HClO3(aq)
 Hydrogen chlorate
becomes
---ic acid
Chloric acid
◦ H3BO3(aq)
 Hydrogen borate
Boric acid
◦ HNO3(aq)
 Hydrogen nitrate
Nitric acid
Lets try some!

RULE 3: hydrogen ---ite
◦ HNO2(aq)
 Hydrogen nitrite
becomes
---ous acid
Nitrous acid
◦ H2ClO2(aq)
 Hydrogen chlorite
Chlorous acid
◦ H2SO3 (aq)
 Hydrogen sulfite
Sulfurous acid