IGCSE Coordinated Sciences Y1
Unit 2: Chemistry
C01, C03, C8.4, C4
Key Notes

• Read pages 47 in your textbook.
• What is a chemical change?
• How is it different from a physical change?
• Describe some indicators that tell us a chemical change has occurred.

Demo
Heating Wax
Burning a Candle
Sugar and Water
CuCl
2 and Al foil
Tearing Paper
Colored Water
KI and Pb(NO
3
)
2
Vinegar & Baking
Soda
Physical? Chemical? Both?
Explanation

Evidence of Physical Changes
• Change in phase or state
• Example: liquid → solid
• Mass or volume change
• Color change that is not permanent

Evidence of Chemical Changes
• Energy is given off or taken in
• Light is observed
• Heat is given off
• Bubbling or fizzing
• Permanent color change

• Important vocabulary:
– atom,
– element,
– molecule,
– compound,
– mixtures
• Write a definition & draw a picture for each term using pg. 30, 31, 46, & 54
• Then, write about the similarities and differences between each.

• An element is unique because it…
• Molecules are similar to compounds because…
• A molecule is different than a compound because …
• A mixture is different than either a compound or molecule because…

• Read & take notes on the description of an atom on Page 32 - 33.
• Watch the video clip about structure of an atom.

• Fill in the table with information about the different parts of an atom:
Charge Location in atom
Relative
Mass
Particle
Proton
Neutron
Electron

• We are going to play a game to determine what you remember about atoms.
• As you start to recognize patterns, fill in your chart. Have your teacher check your work, then if asked, write up an answer on the group chart.
• Discover the patterns and then write the rules if you figure them out!

• Atomic number
• Atomic mass / nucleon number
• Periods vs. Families
• Metals, transition metals, non-metals, noble gasses

e Element p +
Hydrogen
Sulfur
Iodine
Sodium n
0

• How can an element like Chlorine have an atomic mass of 35.5?
• Isotopes are atoms of the same element
(same proton number) with different numbers of neutrons.
• Most elements have isotopes!

• Chlorine – 35 has an atomic mass of 35.
How many protons does it have?
How many neutrons?
• Chlorine – 37 has an atomic mass of __.
p + = _____ n o = ______
• Of all the chlorine atoms in the world, 75% are chlorine-35 and 25% are chlorine-37, giving an average atomic mass of 35.5.

• Electrons travel around the nucleus in orbitals.
• We will study the first 3 orbitals:
– The first orbital can hold up to 2 electrons
– The second orbital can hold up to 8 electrons
– The third orbital can hold up to 8 electrons
• Orbitals always fill from the inside out.
• Example: Lithium

• Example: Lithium

Try these:
• Hydrogen
• Aluminum

• Obtain a “building atoms” sheet on Haiku.
• Use your baggie of particles to build each atom from the sheet, recording your results in your journal.
– Be sure to put the correct number of electrons in each orbital and show the number of protons and neutrons in the nucleus!
– Put your name on your baggie and save your particles for later. You’ll get to eat them later 

• Scientists use a shorthand to denote how many electrons are in each orbital or energy level of an atom.
• See Page 36 for an example!
• Go back to your atom drawings and write the electronic configurations for each atom you built.

A FULL outer shell
(valence electrons)

ION: an atom that has gained or lost electrons.
• Cation: An atom that has lost one or more electrons. Positively charged.
• Anion: An atom that has gained one or more electron. Negatively charged.

How many electrons must be gained or lost by each atom?
• Lithium
• Argon
• Chlorine
• Potassium
• Phosphorus
• Magnesium

• Group 1 loses one electron; they all have +1 charge.
• Group 6 gains two electrons; they all have a
–2 charge.
– Group I
– Group II
– Group III +3
– Group V
– Group VI
-3
– Group VII -1
– Group VIII 0
+1
+2
-2

• How do metals and nonmetals come together to form IONIC bonds?
– Explain in your notes using an example from your reading on page 50 – 51.
– Show using your baggie of candy “particles” the formation of:
• Lithium + Fluorine
(how many of each atom do you need?)
• Lithium + Oxygen
(how many of each atom do you need?)

• When sodium and chlorine bond, they form a lattice structure made up of repeating units of positive and negative charges.
• Building a lattice structure

• Determine the formula for the compound made when the following atoms combine…
– Magnesium + Bromine
– Sodium + Oxygen
– Aluminum + Fluorine

• How to write and name ionic compounds!
– When writing the formula, first write the symbol and charge of each ion.
• Switch the charges and write them as subscripts.
• Simplify!
– When naming a compound, write the name of the cation first then the anion with an –ide ending.
– Use roman numerals for ions with more than one charge [i.e. Fe +3 would be Iron (III), Fe +2 is Iron
(II)]

Sodium, Na +1
Chloride, Cl -
1
Bromide,
Br -1
Oxide, O -2
Magnesium,
Mg +2
Aluminum,
Al +3

Sodium,
Na +1
Magnesium,
Mg +2
Aluminum,
Al +3
Chloride,
Cl -1
NaCl
MgCl
AlCl
3
2
Bromide,
Br -1
NaBr
Oxide, O -2
Na
2
O
MgBr
AlBr
3
2
MgO
Al
2
O
3

• A Polyatomic Ion is an ion made up of more than one type of atom.
• Most have a negative charge, with the exception of ammonium.
• Ionic Compounds with polyatomic ions are written and named in the same manner as binary compounds.

The only exceptions are:
• Each polyatomic ion is treated as one element
• Parenthesis are used around the entire ion when expressing the quantity necessary to balance out the charge
– Ex. Cu +2 + OH -1  Cu(OH)
2
Hydroxide and is named Copper (II)
• Polyatomic ions are named according to their own naming convention
– Ex. Li
3
PO
4 is called Lithium Phosphate

• Ca +2 + PO
4
-3
• Mg + ClO
4
• Cu +1 + SO
4
• Aluminum Hydroxide
• Iron (II) + Permanganate

Sodium,
Na +1
Magnesium
, Mg +2
Aluminum,
Al +3
Ammonium
, NH
4
+1
Hydroxide,
OH -1
Nitrate,
Carbonate,
CO
3
-2
Sulfate,
SO
4
-2

Sodium,
Na +1
Magnesium
, Mg +2
Hydroxide,
OH -1
NaOH
Nitrate,
NO
3
-1
NaNO
3
Carbonate,
CO
3
-2
Na
2
CO
3
Sulfate,
SO
4
-2
Na
2
SO
4
Mg(OH)
2
Mg(NO
3
)
2
MgCO
3
MgSO
4
Al(OH)
3
Al(NO
3
)
3
Al
2
(CO
3
)
3
Al
2
(SO
4
)
3
Aluminum,
Al +3
Ammonium
, NH
4
+1
NH
4
OH NH
4
NO
3
(NH
4
)
2
CO
3
(NH
4
)
2
SO
4

• 1. Do the “Polyatomic Naming” sheet found on Haiku.
• 2. IF more time remains, do another practice sheet from “Ionic Bonding” block on Haiku

• Covalent bonds occur between two or more non-metals .
• Valence electrons are shared to complete atoms outer shells.

Review:
Ionic bonds occur between metals and nonmetals.

• CO
2
• CO
• H
2
O

• Dot and Cross Structures & Line structures
• Hydrogen + Hydrogen
• Hydrogen + Oxygen

• Bromine + Bromine
• Oxygen + Oxygen

• Carbon + Oxygen

• Build each covalent compound on the half sheet.
• Draw a color coded 3-D picture of the actual compound in your notebook.
• Draw a dot and cross structure for each compound beneath your drawing.

• There should be no empty holes in your atoms
• There should be no empty bonds in your compounds
• You cannot build rings of 3 or 4 atoms

• Carbon = Black
• Hydrogen = White
• Oxygen = Red
• Halogens = Green
• Nitrogen = Blue (ignore one hole)
• Bonds = Gray Tubes

• Describe how a molecule of ammonia
(NH
3
) is made. Include a description of how each covalent bond is formed.

• To finish out our study of Covalent
Compounds, read about:
– Giant Covalent Structures (pg. 60 - 61)
– Simple Molecular Structures (pg. 56 - 57)
• How does the structure of a compound or molecule affect its properties?

• What are the differences in the properties of ionic vs. covalent compounds?
• Read pg. 58 - 59

Work on these problems as I check off your homework:
What is wrong with these formulas?
•
•
• Zn +2 + O -2

Zn
2
O
2
NH
Fe
4
+
+3
+ NO
+ CO
3
-
3
-2

NH
3
NO
4

Fe
2
CO
33

• Reactants  products
Na
2
SO
4 (aq)
+ CaCl
2 (aq)
 CaSO
4 (s)
+ 2NaCl
(aq)
• The state of each reactant and product can be shown with a small symbol in parenthesis at the end of each formula (see above).
– s for solid, l for liquid, g for gas, and aq if the compound is an aqueous solution

• Lead (II) nitrate solution reacts with potassium iodide solution to form lead (II) iodide solid and potassium nitrate solution.
• Copper metal reacts with oxygen gas (O
2
) to form copper (II) oxide.

• Iron metal reacts with water to form iron
(III) oxide solid and hydrogen gas (H
2
).
• magnesium hydroxide solution + carbon dioxide  water + solid magnesium carbonate

• When a chemical reaction happens bonds are broken, elements get rearranged, and new compounds are formed.
• The Law of Conservation of Matter tells us matter (or mass) cannot be created or destroyed!
– The number of atoms of each type is the same in both the reactants and products!

• Baking Soda & Vinegar

• Step 1: Take an atom inventory
• *List the type and number of each atom on both sides of the chemical equation
• H
2
CO
3
• Ba(OH)
2
• C
3
H
8
• NaC
2
H
3
O
2
• Al(NO
2
)
3

• Step 2: Law of Conservation of Mass
*We must have the same number of atoms of each type on both sides
*Use Coefficients to balance
~Example: 2NH
3
H
2
+ O
2

H
2
O

• FeS + HCl

FeCl
2
+ H
2
S
• Fe
2
O
3
+ H
2

Fe + H
2
O

• P + O
2

P
2
O
5

• Work through the practice problems on the
Balancing Equations practice sheet.
(Download from Haiku).
– Show your atom inventory for every problem!
• Do the “Writing and Balancing Equations
WS” (Find on Haiku).