Science 10 Review - Chemistry

What is chemistry?

 the study of changes in matter
Classifying matter
Matter
Heterogeneous
(Non-uniform mixtures)
Homogeneous
(Uniform matter)
Pure substances
Compounds
Elements

matter: anything that has mass and occupies space.

heterogeneous substances: non-uniform and may consist of more than one phase (e.g. Coca-Cola, ice
cream float)
homogeneous substances: uniform and consist of only one phase (e.g. tap water)
 homogeneous mixtures (solutions) can be separated by physical means such as filtering (e.g.
Kool-aid)
pure substances can’t be separated by physical means (e.g. water)
 compounds can be separated by means of a chemical change into its different atoms (e.g.
water can be separated into hydrogen and oxygen)
 elements cannot be broken down into simpler chemical substances



Homogeneous
mixtures
(Solutions)
The periodic table

based on rules established
by IUPAC (International
Union of Pure and Applied
Chemistry), elements are
displayed in a chart called
the periodic table
 the characteristics
of the elements
listed on the
periodic table are
accurate at SATP
(standard ambient
temperature (25°C)


and pressure (100kPa))
Split into three categories:
 Metals, non-metal and metalloids
 Each one of these groups contains elements with similar chemical and physical properties.
features of the modern periodic table:
 family / group
 vertical column with similar chemical properties
 alkali metals: (1) soft, silver-coloured, react violently with water to form basic solutions
 alkaline-earth metals: (2) light, reactive metals that form oxide coatings when exposed to air
 transition elements: exhibit a wide range of chemical and physical properties
 halogens: (17) extremely reactive non-metals
 noble gases: (18) extremely low chemical reactivity
 period
 horizontal row of elements whose properties gradually change from metallic to nonmetallic
as you move from left to right



metals / non-metals
 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
o Reactive: such as sodium, which will combust into a flurry of flame when it come on
contact with air.
o Inert: extremely un-reative, like platinum or gold
Non-metals
 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.

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 number (= number of protons)
atomic mass (mass of one mole of that substance.
1 mole = 6.02 x 1023 molecules)
Symbol/ Name
common/other ion charge (how many electrons
gained/lost by atoms to form ions

Atomic Theory
atoms:




based on these laws, John Dalton introduced a new atomic theory of matter in 1803.
 all matter is composed of tiny, indivisible particles called atoms
 all atoms of the same element have identical properties, while atoms of different elements have
different properties
 atoms of two or more elements can combine in constant ratios to form new substances
electrons:
 discovered in the late 1800s by J.J. Thompson
 tiny, negatively charged particles called electrons could be separated from atoms; that is, that
atoms could be further divided into smaller particles.
 the net charge of the atoms was zero (neutral)
nucleus & protons:
 shortly thereafter, Ernest Rutherford revised the previous models:
 atoms contain a tiny, positively-charged core called the nucleus, which is surrounded by mostly
empty space containing negative electrons. Though tiny, the nucleus makes up most of the
mass of the atom.
 the nucleus is made up of tiny particles called protons. Each element has a different number of
protons ( its atomic number).
 later, another scientist also found the nucleus contained equally small, but neutral, particles
called neutrons
electrons & orbitals:
 Niels Bohr discovered that electrons occupy fixed orbits, like planets orbiting around the sun.
 electrons cannot exist between levels, but they can move from one level to another.
 each level has a fixed capacity for electrons
 the shell closest to the nucleus can only hold 2 electrons,
 the second and third can both hold 8
 if a shell is full, new electrons cannot move to that shell
 each shell represents an amount of energy held by that electron
 electrons further from the nucleus have more energy than those in the closest orbit.
 an element with its outermost orbit full is stable and unreactive
particle
proton
neutron
electron
symbol
p+
n
e-
charge
1+
0
1-
mass
1.7 x 10-24-g
1.7 x 10-24-g
9.1x10-28g
notes
make up nearly all the mass of an
atom, but hardly any volume
barely any mass, but most of the
volume
location
in the nucleus
around the nucleus

Summary of what we know about the atom
 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.
 To make it simpler we use Mass Number.
 protons + neutrons = mass number
 Therefore…
 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

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.
 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.
~ Anions: pick up (gain) electrons. Since they now have more electrons than protons, they have a
negative charge.
 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.
Lets try some! What would the following elements be as ions?
Magnesium, nitrogen, 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.
~ The number of electrons that can fill each orbital is equal to the
number of elements in each row.
~ The number of valance electrons is equal to the number of the column that the element is in.

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.
~ Draw in the correct number of orbital

Recall the number of rings = column the
element is in

The electrons are placed on the rings


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.
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
~ 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.
 How to write Ionic Cmpds:
~ Step 1: Write each ion with its charge
 Remember: Certain columns on the periodic table are always the same, and some
have more than one choice.

Ex. calcium and bromine: Ca2+ and Br –
~ Step 2: Figure out how many of each you need to make the charges balance.
 The best way is to do a switch-er-oo between the number of atoms and the opposite
ion’s charge.
~ Step 3: Write the formula using subscripts to show how many of each atom you need AND
the state


CaBr2 (s)
Note: Ionic compounds are always solid (s)
 Multiple charges…
~ These are called multivalent ions.
~ In order to know which of the ion charges you need the formula will have a roman numeral in it.
 Cr2+ is written chromium (II)
 Co3+ is written cobalt (III)
 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.
~ 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.


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.
Let’s try some!
RULE 1: hydrogen ---ide
becomes
hydro ---ic acid
HF(aq)
H2P(aq)
HI (aq)
RULE 2: hydrogen ---ate
becomes
---ic acid
HClO3(aq)
H3BO3(aq)
HNO3(aq)
RULE 3: hydrogen ---ite
becomes
---ous acid
HNO2(aq)
H2ClO2(aq)
H2SO3 (aq)
 Solubility:
~ the property of a solid, liquid, or gaseous
chemical substance called solute to dissolve
in a liquid solvent to form a homogeneous
solution.
~ The solubility of a substance depends on
the specific solvent as well as the
temperature and pressure
~ To predict the solubility of ionic compounds
only we use a solubility table.
~ The most up-to-date version of the solubility
table is on page two of your data-booklet.
~ Things that are “soluble” will dissolve and
form a solution.
~ Things that are “slightly soluble” will NOT
dissolve but form a precipitate (chunks that float
viable within the water)
Lets Practice!
Li2SO4,
CaSO4,
NH4F,
LiF,
Rb2O,
Cs3As,
 Chemical Rxns
~ There are 5 types of Chemical reactions:
~ Formation: creating a compound from two elements
 element + element → compound
 Na(s) + Cl(g) →NaCl(s)
~ Simple decomposition: this is when a compounds breaks into its components
 compound → element + element/polyatomic ion
 H2SO4(aq) → H+ (aq) + SO4 2-(aq)
AgCl
~ Single replacement: (one element replaces an element in the compound to make a new
element/compound combo)
 element + compound → new element + new compound
 H2SO4 (aq) Zn(s) → H2(g) + ZnSO4(s)
~ Double replacement: (all the ions switch partners)
 compound + compound→ new compound + new compound
 HCl(aq) + NaOH(aq) → NaCl(s) + HOH(aq)
~ Combustion: (the burning of a hydrocarbon in the presence of O2)
 hydrocarbon + O2 → CO2 + H2O
 C6H12(g) + O2(g) → CO2(g) + H2O(l)
 Rules:
1. In all chemical reactions: atoms, mass and energy are conserved
(There is the same amount of each on either side of the arrow)
2. Balance the chemical reaction by:
a. Write the correct chemical formulas for the reactants and products
b. Identify all the atoms on the reactant side and identify all the similar atoms on the product
side.
c. Compare the numbers of atoms on each side and try to balance them by adding coefficients
on each side.
Lets Practice!
Na (s) + Br2 (l)→
Mg (s) + F2 (g) →
Calcium and oxygen produces?
Magnesium phosphide →
Fe(s) + Cu(NO3)2(aq) →
Zinc and iron (II) chloride →
CaO(s) →
C6H12O6(s) + O2(g) →
Copper (II) nitrate + potassium bromide →

Molar Mass
~ To calculate the molar mass of a compound, add the individual
atomic molar masses of all the atoms in the compound together.
Let’s Practice!
~ Calculate the molar masses of the following compounds:


Water:

Glucose (C6H12O6):
Molar Mass Calculations
~ Manipulate the formula to solve for mass
~ How many moles are in 88.02g of carbon dioxide?
~ What is the mass of 4.38 moles of carbon dioxide?
Physics Review

What is Physics?
 The branch of science concerned with the nature and properties of matter and energy.

Thermodynamics
 The First law of thermodynamics a re-statement of
the Law of Conservation of Energy

Energy cannot be created or destroyed; it can only
be converted from one form to another

EXCEPT in the Law of Thermo-dynamics one of the forms of energy is always heat.

Second Law of Thermodynamics piggy- backs on the
1st Law agreeing that the quantity of energy remains
the same in a system…

HOWEVER the usable quality of the energy
deteriorates over time.
 Each time energy is converted from one
form to another, some of the useful energy is
always “lost” (reduced to a lower-quality, less
useful form).
 The 2nd Law expresses reality: no system can
convert energy from one form to another
with 100% efficiency.
Most often, this “lost” less useful energy in the form
of thermal energy; heat as a result of friction and movement within the system.


The amount of “useful output” relative to the amount of energy input is referred to as the
machines efficiency.

Can be found using the formula:
percent efficiency =

total work output
 100%
total work input
Rearrange the formula for percent efficiency to solve for:
~ total work output
~ total work input
Lets Practice!

A Bunsen burner supplies 4.00 x 103 J of heat to a small beaker of water. Only 125 J of heat is
gained by the beaker and water. Calculate the percent efficiency of the burner.

A small electric motor has an efficiency of 85%. In lifting a small load, it produces 15 J of
mechanical energy input. Calculate the useful mechanical energy output of the motor.
Systems
 Recall the three different conditions that energy transfers can occur in:

Open system: a system that exchanges matter as well as energy with the surroundings.


Closed system: A system that can exchange energy but NOT matter with its surroundings.
Isolated system: A system that can exchange neither energy nor matter with its surroundings.
Truly isolated physical systems do not exist in reality.
Significant Figures/ Digits
 The number of significant digits in an answer to a calculation will depend on the number of
significant digits in the given data:
When are Digits Significant?
 Non-zero digits are always significant.
 22 has two significant digits (2 and 2).
 22.3 has three significant digits (2, 2 and 3).

 With zeroes, the situation is more complicated:
Zeroes placed before other digits are
not significant;
 0.046 has two significant digits
(4 and 6, b/c the preceding zeros DO NOT count).

Zeroes placed between other digits are always significant;
 4009 kg has four significant digits (4, 0,0 and 9).

Zeroes placed after other digits but behind a decimal point are significant;
 7.90 has three significant digits (7, 9 and 0).
Why?
 Values to the right of a decimal place can be rounded as necessary, values to the left of a decimal
CAN NOT.
 So that we can round appropriately to give an answer with the correct number of digits,
scientific notation is used.
 If the answer comes to 8219.10
 8.219 x 103 has four significant digits
 8.22 x 103 has three significant digits.
 8.2 x 103 has two significant digits
Sig Dig’s for Multiplication & Division Functions
 The number of significant digits in an answer should equal the least number of significant digits in
any one of the numbers in the question
 Ex.
What is the speed of a jet plane that travels 528 meters in 4.0 seconds?
v = ∆d
∆t
v = 528m
*528 has 3 sig digs (5, 2 and 8), 4.0 has two
4.0 s
(4 and the trailing 0). So your answer can
v = 132
have only 2 digits *
v = 1.3 x 10 2 m/s
Sig Dig’s for Addition & Subtraction
 When adding and subtracting, your answer must have the same number of decimals places as the
least number of decimal places (not total number of digits) in the in the question
 Ex:
5.67 J
1.1 J
* 5.67 J has two decimal places,
+ 0.9378 J
1.1 J has one decimal place,
7.7078
0.9378 J has four decimal places
* Your answer can have only one decimal
place*
Ans = 7.7 J
Intermediate Answers
 ALWAYS Keep AT LEAST One Extra Digit in Intermediate Answers
 When doing multi-step calculations, keep at least one more significant digit in intermediate
results than needed in your final answer.
 If the final answer requires two significant digits, carry at least three significant digits in
calculations (or just keep them in your calculator).
Scalar
Vector
◦To posses only magnitude, (numerical value only)
◦Ex. It is 10 meters to the neighbors.
Distance
◦The amount of space between two things.
◦Ex. It is 389 km from Edmonton to Jasper
Speed
◦The rate of travel.
◦Ex. We travelled at 110km/h to get from Edmonton to Jasper.
Time
◦The method of measuring the passage of time.
◦Ex. Science class is 60 min long.
◦Ex. It takes 3 ½ hours to get to Jasper
◦To have both magnitude, (numerical value) AND direction.
◦A vector symbol (v ) is used above the value.
◦Ex. It is 10 meters EAST to the neighbors.
Displacement
◦The amount of space between two things IN A GIVEN
DIRECTION.
◦Ex. It is 389 km[West] from Edmonton to Jasper
Velocity
◦The rate of travel IN A GIVEN DIRECTION.
◦Ex. We travelled at 110km/h [West] to get from Edmonton to
Jasper
Acceleration
◦The act of increasing in speed (or velocity) over a period of
time.
◦Ex. You accelerate from 60km/h to 100km/h to merge onto
the highway
Formula Manipulation
 Rearrange the formula to solve for:
 Δd
~ Δv
~ Δt
~ Δt
~ vf
~ vi
~ Δt
~ vi
~ vf
~ vi
~a
 What is Biology?
 The branch of science concerned with the science of living matter in all its forms especially with
reference to origin, growth, reproduction, structure, and behavior.
 Biosphere
 Earth's zone of air, soil, and water that is capable of supporting life, traditionally thought of as a zone
reaching about 10 km into the atmosphere and down to the deepest ocean floor.
The envelope of gases
surrounding the
Earth
The part of the Earth
composed of water including
clouds, oceans, seas, ice caps,
glaciers, lakes, rivers,
underground water supplies,
and atmospheric water vapor.
 Biomes
 A biome is a large geographical region with a
specific climate that the plants and animals
that inhabit it are adapted to.
 Cells and biomes are open systems that
exchange matter and energy with their
surroundings.
 Biomes can be broken down into specific
ecosystems.
 a system formed by the interaction
of a community of organisms with their environment.
The solid outer layer of
the Earth; includes both
land area and the crust
beneath the oceans and
other water bodies.
 Ecosystems consist of biotic & biotic factors
Biotic Factors
 living organisms in the environment
 examples:
 animals
 plants

Abiotic factors
physical, non-living parts of the environment
~ examples:
 water supply
 light
 soil quality
 climate / temperature
 Life Processes’
Animals
Must eat other things in order to meet their nutrition
needs.
Plants
Plant can produce their own food from the sun.
 Some animals eat only plants (herbivores)
 Some animals eat other animals (carnivores)
 Some eat both (omnivores)
When animals ingest nutrients it must first be broken
down physically (by chewing and digesting) and
chemically (by the acids/ enzymes) in the stomach and
digestive tract.
Instead of ingesting nutrients, plants are able to
combine elements in it’s environment to produce food.
Cellular Respiration
Once food has been broken down into small enough
particles, it is taken into the blood and distributed to
each cell.

At the cellular level digestion is referred to as
cellular respiration
The series of reactions and processes that take place in
the cells of organisms to convert food into energy (ATP)
Photosynthesis
The process of converting light energy to chemical
energy and storing it in the bonds of sugar.

Plants only need light energy, CO2, and H2O to
make sugar.
The process of photosynthesis takes place in the
chloroplasts, specifically using chlorophyll, the green
pigment involved in photosynthesis.
Cellular respiration is a process that requires oxygen and
nutrients and gives off carbon dioxide and water and
energy for the organism.
C6H12O6 + 6O2 → 6CO2 + 6H2O + Energy
6CO2 + 6H2O + light energy → C6H12O6 + 6O2