Chemistry Notes - Grade 10
5.1 Properties and Changes
Physical Change: Water to Ice
Chemical Change: Hydrogen peroxide to Water
Physical Property:
⤷ Characteristic/description of the substance
⤷ Based on Qualitative Observations (senses)
⤷ Also based on Quantitative Observations (calculations and measurements)
(e.g colour, texture, density, smell)
Chemical Property:
⤷ Describes how a substance behaves in the presence of another substance (e.g. how it
reacts with other elements and compounds)
⤷ Examples are flammability, corrosion, bleaching, rusting etc.
Physical Change:
⤷ No new substance created
⤷ Most can be reversed
⤷ E.g. Change of state, dissolving sugar in water, cutting paper
Chemical Change:
⤷ Produces a new substance
⤷ Most can’t be reversed
⤷ E.g. Forest fire, boiling an egg, burning wood
Five Clues that a Chemical Change has Occurred:
1. A new colour is observed
2. A new odour is produced
3. Heat or light is produced or absorbed
4. Bubbles of gas are formed
5. A solid material (precipitate) forms in a liquid
5.4 Patterns in the Periodic Table
Where are the metals, non metals, and metalloids located on the periodic table?
⤷ Metals: Left side
⤷ Metalloid: On the staircase
⤷ Non Metals: Right side exception being Hydrogen
Where are the alkali metals found? What are their properties?
⤷ Location: Group 1
⤷ Properties: Highly Reactive, shiny, soft, low melting point, and they have one valence
electrons, soft, ductile
Where are the alkaline earth metals found? What are their properties?
⤷ Location: Group 2
⤷ Properties: Reactive, shiny, silvery, readily found in nature, and they have two valence
electrons
Where are the halogens found? What are their properties?
⤷ Location: Group 17
⤷ Properties: Highly Reactive, reacts well with alkali metals, and they have seven valence
electrons, toxic
Where are the noble gases found? What are their properties?
⤷ Location: Group 18
⤷ Properties: Very stable, odorless, full valence shell, non conductive, glows when in contact
with electricity, standard temperature and pressure
5.5 Atoms & Ions
What are Ions?
⤷ Atoms have an overall charge of zero as they have equal numbers of protons and
electrons
⤷ Ions are formed when an atom or group of them gain/lose at least one electron
⤷ Ions have unequal numbers of protons and electrons, so they have a positive or
negative ionic charge
Positive Ions
⤷ Cation = Positively charged ion
⤷ E.g. A sodium ion loses its outermost electron to form an ion with a stable octet outer
shell
⤷ The name of a cation is the same as the name of the element with “ion” added
(Sodium ion)
⤷ Positive ions are usually from atoms of metal elements
Negative Ions
⤷ Anion = Negatively charged ion
⤷ E.g. The fluorine atom gains an electron in its outermost shell to form an ion with a
stable octet outer
Shell
⤷ To name an anion, add “ide” to the stem of the element’s name (fluoride ion)
⤷ Negative ions are usually from atoms of non-metal elements
Why are Ions formed?
⤷ Atoms gain/lose enough electrons to have a full outer shell of eight electrons, this is
the most stable and less likely to undergo a chemical change
Why are Ions important?
⤷ Many ions are critical for good health, like sodium ions, phosphorus ions, iron ions,
and calcium ions (there’s more)
⤷ However it is critical their concentrations are regulated; too much sodium leads to
heart/kidney disease, and not enough iron leads to anemia
5.9 Polyatomic Ions
What are Polyatomic Ions?
⤷ Groups of atoms that tend to stay together and carry an overall charge
<== NO3
<==SO4
⤷ These behave exactly like simple ions to produce compounds with an overall charge of
0
Naming Ionic Compounds:
⤷ Write the name of the first element (positively charged ion)
⤷ Write the name of the polyatomic ion
Note: If the polyatomic ion happens to be ammonium it is written first because it is the
positive ion
Positive Ion Always Comes First
Examples:
⤷ Na2SO4 = Sodium Sulfate
⤷ K-2CO3 = Potassium Carbonate
⤷ NH2NO3 = Ammonium Nitrate
⤷ Fe3(PO4)2 = Fe=3(2)=6 ; PO=2(-3)=-6 ; 6+(-6)=0
“Iron (II) Phosphate”
⤷ Al(OH)3 = Aluminum hydroxide
⤷ Cu2(CO)2 = Cu=2(2)=4 ; CO=2(-2)=-4 ; 4+(-4)=0
“Copper (II) Carbonate”
Writing Formulas for Polyatomic Compounds:
⤷ Write the symbol of the metal and the polyatomic group
⤷ Place brackets around the polyatomic group
⤷ Write the ionic charge for the metal and polyatomic group
⤷ Cross over charges to subscripts
⤷ Simplify if necessary, but do not reduce the polyatomic group inside the brackets
Examples:
⤷ Sodium Sulfate = Na+(SO­­4)2- = Na2(SO4)
⤷ Calcium Hydroxide = Ca2(OH)- = Ca(OH)2
⤷
⤷
⤷
⤷
Sodium Carbonate = Na+(CO3)2- = Na2(CO3)
Magnesium Nitrate = Mg2(NO3)- = Mg(NO3)2
Sodium Hydroxide = Na+(OH)- = Na(OH)
Copper (II) Sulfate = Cu2(SO4)2 = Cu(SO4) or Cu2(SO4)2
5.10 Molecular Compounds
Molecules and Covalent Bonding
⤷ A molecular compound is a pure substance formed from non-metals.
⤷ Involves the sharing of electrons between the non-metals in order to obtain full
electron orbits.
⤷ Examples include: H2O; CO2; H2
Covalent Bond
⤷ A shared pair of electrons between two nonmetals that hold the atoms together
Types of Compounds
Diatomic Molecules
⤷ Molecules made up of two identical atoms
⤷ Some atoms naturally form diatomic molecules
Examples: H2; O2; Br2; N2; F2; Cl2; I2
⤷ These are all gases – hydrogen gas (H2); nitrogen gas (N2); oxygen gas (O2)
MEMORIZE THESE!!
Naming Molecular Compounds
⤷ When naming molecular compounds, scientists use prefixes to specify the number of
each type of element in a molecule.
⤷ The prefix mono is used only for the second element in the compound.
1.
Write the name of the first elements and change the ending of the second element to
“ide”
2.
Add the prefixes to each element.
(NOTE: The prefix mono- is not necessary on the first element.)
Chemical Formulas for Molecular Compounds
⤷ Write the symbols for each element in the compound.
⤷ Use the prefixes for each element to determine how many of each element is found in
the molecule and add them to the symbols using subscripts.
(NOTE: There will be no prefix in the first element if there is only one.)
Common Names
⤷ Some molecular compounds are referred to
⤷ by their common names.
Examples:
⤷ H2O – water
⤷ NH3 – ammonia
⤷ H2O2 – hydrogen peroxide
⤷ CH4 – methane
⤷ O3 – ozone
⤷ NO – nitric oxide
⤷ HCl – hydrochloric acid
6.3 Conserving Mass in Chemical Reactions
Conserving Mass
⤷ Chemical reactions cause a rearrangement of atoms in the reactants to form different
products
⤷ Molecules might be broken apart and new ones might be put together, but the atoms
in the product(s) are the same as the ones that were present in the reactants
⤷ Notice that all the atoms present at the beginning of the chemical reaction are present
after the reaction. The total mass of the atoms in the reactants remains equal to the total
mass of the atoms in the products.
The Law of Conservation of Mass
⤷ The total mass of the reactants is equal to the total mass of the products!
Reactant A + Reactant B → Product C + Product D
45 g
25 g
50 g
?
Calcium + Hydrochloric acid → Calcium chloride + Hydrogen gas
16 g
33 g
28 g
?
How to Count Atoms
⤷ The symbol of an element represents one atom of that element.
⤷ A subscript indicates more than one atom of the element.
⤷ A subscript outside a bracket multiplies all the elements inside the brackets.
⤷A coefficient is a number written in front of a chemical symbol and indicates the
number of atoms of that element or compound.
⤷A coefficient written in front of a chemical formula multiplies the number of atoms of
each element in the formula
Balancing Chemical Equations
⤷ Word Equation
Methane + Oxygen gas → Water + Carbon dioxide
⤷ Skeleton Equation
CH4 + O2 H2O + CO2
⤷ The chemical equation must obey the law of conservation of mass. Thus the total
number of one type of atom must be the same on both sides of the equation.
⤷ To ensure this one balances the equation by putting COEFFICIENTS in front of the
substances.
~Balanced Chemical Equation~
CH4 (g) + 2 O2 (g) 2 H2O (g) + CO2 (g)
6.5 & 6.6 Types Of Chemical Reactions
Synthesis
⤷ The combination of smaller atoms and/or molecules into larger molecules
⤷ Also called combination or addition reactions The general pattern for a synthesis
reaction is:
⤷ They have two or more reactants and usually one product
⤷Example: 4 Al (s) + 3 O2 (g) 2 Al2O3 (s)
Decomposition
⤷ The breaking down of a large compound into smaller compounds or elements
⤷ The general pattern for a decomposition reaction is:
⤷They have one reactant and two or more products
⤷Example: 2 NaN3 (s) 2 Na (s) + 3 N2 (g)
Single Displacement
⤷ Involves an element and a compound
⤷ Occurs when one element displaces or replaces another element in a compound,
producing a new compound and a new element
⤷ The general pattern for a single displacement reaction is:
⤷ A metal always replaces another metal and a non-metal always replaces another non-
metal
⤷ Example: Cu (s) + 2 AgNO3 (aq) Cu(NO3)2 (aq) + 2 Ag
Double Displacement
⤷ Involves two compounds
⤷ Occurs when two elements in different compounds trade places.
⤷ Two solutions that undergo a double displacement reaction produce an insoluble
compound called a precipitate.
⤷ The general pattern for a double displacement reaction is:
⤷ A metal always replace another metal and a nonmetal always replaces another nonmetal
⤷ Example: Pb(NO3)2 (aq) + 2 KI (aq) PbI2 (s) + 2 KNO3 (aq)
6.9 Combustion Notes
Combustion Reaction
⤷ Involves the burning of hydrocarbons (Fuel) in the presence of oxygen
⤷ Combustion doesn't always result in fire, but when it does, a flame is a characteristic
indicator of the reaction.
⤷ Products are usually an oxide and energy
Complete Combustion
⤷ Hydrocarbons burn completely by releasing the energy they contain
⤷ Occurs when the supply of oxygen is plentiful
⤷The product of complete reaction combustion are always carbon dioxide and water
Example: CH4 + O2 ---> CO2 + H2O + Energy
Incomplete Combustion
⤷ Hydrocarbons are not burned completely, releasing carbon monoxide gas and carbon,
in addition to carbon dioxide and water.
⤷ Occurs when oxygen supply is limited
⤷ Often indicated by an orange flickering flame
Example: CH4 + O2 ---> H2O + Co2 + CO + C + Energy
Other Combustion Reactions
⤷Other substances besides hydrocarbons undergo combustion reactions
⤷Their elements react with oxygen to form oxides
7.2 Acids And Bases
What Are Acids And Bases
⤷Acids and bases are substances that form ions when dissolved in water.
Properties Of Acids And Bases
Acids
Bases
Reacts with metals and
carbonates
WD
Conducts electricity
Conducts electricity
Turns blue litmus paper red
Turns red litmus paper
blue
Tastes sour
Tastes bitter and feels
slippery
pH < 7
pH > 7
Neutralizes bases
Neutralizes acids
Electrical Conductivity
⤷An acid dissolves in water to form hydrogen ions (H+)
Example:
HCl (aq) à H+ (aq) + Cl- (aq)
⤷A base dissolves in water to form hydroxide ions (OH-)
Example:
NaOH (aq) à Na+ (aq) + OH- (aq)
Names And Chemical Formulas Of Acids
⤷There are two common groups of acids:
-Binary Acids
-Oxyacids
Binary Acids
⤷Binary Acids contain only two elements
Acid Name
Chemical
Formula
Hydrofluoric Acid HF (aq)
Hydrochloric
Acid
HCl (aq)
Hydrobromic
Acid
HBr (aq)
Hydrosulfuric
Acid
H2S (aq)
Oxyacids
⤷Oxyacids are related to polyatomic ions
Acid Name
Chemical Formula
Nitric Acid
HNo3 (aq)
Carbonic Acid
H2Co3 (aq)
Sulfuric Acid
H2So4 (aq)
Phosphoric
Acid
H3PO4 (aq)
⤷Many bases are ionic compounds containing hydroxide or carbonate ions.
Base Chemical Formula
Sodium hydroxide NaOH (aq)
Calcium hydroxide Ca(OH)2 (aq)
Magnesium hydroxide Mg(OH)2 (aq)
Sodium bicarbonate NaHCO3 (aq)
⤷Acid-base indicators are substances that changes colour depending on whether it is in
the presence of an acid or a base.
Bromothymol Blue (BTB)
› Acid – yellow
› Base - blue
Acid Base
Phenolphthalein
› Acid – colourless
› Base - pink
Acid Base
Litmus paper
› Acid – red
› Base – blue
Cabbage juice (all pH levels appear as different colours)
Acid- reddish colour
Base- greenish colour