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Classification of Matter. Matter. Mixture. Pure Substance. Homogeneous (solution). Heterogeneous (mechanical mixture). Element. Compound. Classification of Matter (Alternate). Matter. Homogeneous. Heterogeneous. Mixture. Pure Substance. Element. Compound. Physical properties.

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classification of matter
Classification of Matter



Pure Substance

Homogeneous (solution)

Heterogeneous (mechanical mixture)



classification of matter alternate
Classification of Matter (Alternate)





Pure Substance



physical properties
Physical properties
  • A physical property is any aspect of matter that can be measured or seen without changing the composition of the matter.
  • Qualitative
    • Odor, color, texture, state, malleability
  • Quantitative
    • Melting point, boiling point, density, mass
physical change
Physical Change
  • Does not change the composition of the matter - doesn't change what the substance is
  • In a mixture, individual components retain properties of the original mixture; e.g. dissolving salt in water
    • Change of state (melting, freezing, etc.) is an example of a physical change
physical separation
Physical Separation
  • Separating a mixture based on physical properties
  • Involves a physical change only (not chemical)
  • Individual components have properties of the mixture
  • Examples:
    • Filtration
    • Using magnetic properties
    • Sedimentation, using density differences
chemical properties
Chemical Properties
  • A chemical property is any property of matter that becomes evident during a chemical reaction.
  • Can only be measured by changing a substance's chemical identity.
  • Chemical properties cannot be determined just by viewing or touching the matter.
chemical change
Chemical Change
  • A new substance is formed and energy is either given off or absorbed
  • Involves energy
    • If heat is given off during the reaction, than the reaction is considered to be exothermic.
    • If heat is required for the reaction, than the reaction is considered to be endothermic.
  • Composition of the substance is altered
  • New substances are produced with properties different from the original substance
  • Not easily reversed
evidence of chemical change
Evidence of Chemical Change
  • The following may indicate that a chemical change has occurred
    • Colour Change
    • Temperature Change
    • Odour Given Off
    • Precipitate is Formed
    • Gas Produced
    • Any new substance produced
  • Examples
    • Burning
    • Metal in acid
    • Electrolysis of water
chemical separation
Chemical Separation
  • Separating a substance using a chemical change
  • Can be used to separate compounds
the atom
The Atom
  • Three basic components of an atom:
    • Electrons
    • Protons
    • Neutrons
  • An atom is mostly empty space, with almost all of the mass contained in the nucleus
  • Protons and Neutrons are contained in the centre of the atom known as the nucleus.
  • Electrons “orbit” the nucleus.
parts of the atom
Parts of the Atom
  • Neutron: Large with no charge (n0)
  • Proton: Large with a positive charge (p+)
  • Electron: Small with a negative charge (e-)
  • Charge on an electron is equal and opposite to the charge on a proton
  • All elements (in their ground state) are neutral, meaning the number of protons and the number of electrons are equal.
representing elements
Representing Elements
  • To represent elements we use the symbol




  • X – Atomic symbol
  • Z – Atomic number (p+)
  • A – Mass number (p+ + n0)
  • Na  Sodium
  • Atomic Number = 11  11 protons
  • Mass Number = 23  23 - 11 = 12 neutrons
  • Neutral  # electrons = # protons = 11




periodic table
Periodic Table
  • 118(?) Elements are arranged in Groups (columns) and Periods (rows)
  • There are three types of elements:
    • Metals
    • Metalloids
    • Non-Metals
  • Periodic Table is broken into sections for each type
patterns and trends
Patterns and Trends




Elements that border the “staircase” tend to have both metal and non-metal properties. These elements are known as metalloids.

electron configurations
Electron Configurations
  • Electrons “orbit” the nucleus in regions known as shells.
  • The elements in the first period have one shell and each period adds another.
  • The first shell can only hold 2 electrons and each shell after that can hold 8 electrons.
  • When a shell is full move on to the next one.
outer shell electrons
Outer Shell/Electrons
  • The outer most shell of an electron configuration is known as the valence shell.
  • Electrons contained in this shell are known as valence electrons.
  • What do you notice about elements of the same group and the number of electrons in the outer most shell?
period table and configurations
Period Table and Configurations
  • Period determines the number of shells and the group determines the number of valence electrons.
  • When looking at chemical similarities in the periodic table we look at the groups and not the periods.
common names of groups
Common Names of Groups
  • Group 1 (1A) – Alkali Metals
  • Group 2 (2A) – Alkaline Earth Metals
  • Groups 3-12 are called the transition metals
  • Groups 1,2,13-18 are called representative elements
common names of groups1
Common Names of Groups
  • Group 17 (7A) – Halogens
    • Means “salt former”
  • Group 18 (8A) – Noble Gases
    • Only for first six periods.
    • Called Noble Gases because they don’t easily react with the other elements (full valence shell)
  • Isotopes are variations of the element where neutrons are added or removed to give different types (weights) of atoms.
  • Still the same element just different mass number.
  • Atomic number does not change so there is still the same number of protons present.
  • The atomic mass seen on the periodic table is the average mass of all the isotopes of that element.
isotope example hydrogen
Isotope Example (Hydrogen)

0 neutrons 1 neutron 2 neutrons

  • 2H and 3H are known as deuterium and tritium.
  • Also called “heavy” hydrogen
lewis dot diagrams
Lewis Dot Diagrams
  • Lewis Dot Diagrams are used to just represent the valence electrons.
  • Hydrogen:
  • Helium:



how to fill in lewis dot diagrams
How to Fill in Lewis Dot Diagrams
  • Start at the top and then fill in going clockwise.
  • Can only have up to a maximum of 8 dots around the atomic symbol.




charged elements
Charged Elements
  • Elements can become charged if the there is a change in the number of electrons
  • Elements try to get to the electron configuration of the closest noble gas –full valence shell.
  • Charged elements are known as ions.
    • Positively Charged Ion  Cation
    • Negatively Charged Ion  Anion
  • Boron has 3 electrons in its valence shell
    • Must lose 3 electrons to achieve a full valance shell
    • Acquires a charge of +3
    • Ion is B3+
  • How elements react depends on their valence electrons.
positively charged
Positively Charged
  • Generally, metals tend to lose valence electrons relatively easily.
  • Elements that can easily lose an electron are known as electron donors.
  • To remove an electron from the Group 1 metals requires relatively little energy.
  • As you move down the group it becomes easier to remove the valence electron.
  • Metals are less reactive as you move across the period to the right.
negatively charged
Negatively Charged
  • Since non-metals have a greater number of valence electrons, they must gain electrons to fill their valence shell.
  • Elements that can easily gain an electron are known as electron acceptors.
  • The Halogens are very reactive elements.
  • As you move down the group the elements become less reactive.
  • Non-metals are less reactive as you move across the period to the left.
forming compounds
Forming Compounds
  • When two atoms collide, the valence electrons of each atom interact.
  • Elements try to get to the electron configuration of the closest noble gas - full valence shell.
valence shell
Valence Shell
  • Three ways for an atom to acquire a full valence shell.
    • An atom may give up electrons
    • An atom may gain electrons
    • An atom may share electrons
  • When atoms give up and gain electrons in a reaction the resulting compound is known as an ionic compound with an ionic bond.
  • The third way to acquire a full valence shell will be talked about later in the course.
ionic compounds
Ionic Compounds
  • Ionic compounds involve bonds between a metal cation and a non-metal anion.
  • If just two different elements are involved, than you have a binary compound.
  • Binary compounds require that the total charge (sum of the element’s charges) of the compound is equal to zero.
  • We represent the compound by writing down the element symbol for cation first and then the anion
  • Subscripts after each symbol identify how many ions are required for a total charge of zero.
  • The representation of the compound is known as the chemical formula
lewis structures
Lewis Structures
  • Sodium (Na+) and Chlorine (Cl-)
  • Now Chlorine has a full valence shell and the ionic compound NaCl is formed.



[ ]+

[ ]-



  • NaCl (Table Salt)
  • How do we name this compound?
  • Sodium Chloride
  • The suffix “ide” is put at the end of the name for the element that is the electron acceptor (anion)

Na+ + Cl- NaCl

  • The sodium has a +1 charge and the chlorine has a -1 charge therefore +1 + -1 = 0.
another example
Another Example
  • What would happen if we combined Magnesium and Chlorine?
  • Charges do not add up to zero.
  • Therefore we need more of one of the elements, but which one.
  • Magnesium has a 2+ charge and Chlorine has a 1- charge so we need two Chlorine.
  • MgCl2 (Magnesium Chloride)
  • Can also be done by drawing out the required number of atoms to get a total charge of zero.
polyatomic ions
Polyatomic Ions
  • Polyatomic Ions consist of two or more non-metal atoms grouped together.
  • There is only one common polyatomic cation
    • Ammonium NH4+
  • There are several common polyatomic anions
    • Hydroxide OH-
    • Carbonate CO32-
    • Nitrate NO3-
    • Sulfate SO42-
    • Chlorate ClO3-
    • Phosphate PO43-
polyatomic i ons
Polyatomic Ions
  • Compounds are named the same way
  • Writing the chemical formula is a little different - If more than one polyatomic ion is needed, than brackets must be put around the ion
  • Example: The chemical formula for ammonium oxide is (NH4)2O not NH42O
  • Do not forget the brackets!!!!
polyatomic example
Polyatomic Example
  • Calcium Nitrate
  • Calcium (Ca2+) and Nitrate (NO3-)
  • Need two Nitrate ions to balance charges.
  • Ca(NO3)2
transition metals
Transition Metals
  • Transition metals can form more than one ion - except for silver(+1), zinc (+2) and aluminum (+3).
  • For example Sodium can only produce the Na+ ion. Iron on the other hand can produce two ions.

Fe  Fe2+ or Fe3+

  • A roman numeral is placed after the atom in brackets to identify the charge
  • Iron that produces the +2 ion is iron(II)
  • Iron that produces the +3 ion is iron(III)
  • 1. Iron(III) Oxide (Rust)

Fe3+ O2-

Fe3+ O2-


  • Charge of +6 from the iron and -6 from the oxygen. Chemical Formula - Fe2O3
  • 2. CuCl2

Cu Cl-


  • Cu must have a +2 charge to balance the -2 from the 2 Cl. Copper(II) Chloride
covalent bonds
Covalent Bonds
  • Two or more non-metallic elements.
  • Electrons must be shared since both atoms are looking to gain electrons.
  • When atoms share electrons they are joined by a covalent bond.
  • A neutral particle that is composed of atoms joined together by covalent bonds is called a molecule.
  • Substances that are composed of molecules are called molecular compounds.
molecular compounds
Molecular Compounds
  • Water (H2O)
  • Two H+ atoms and a O2- atom.







naming molecular
Naming Molecular
  • H2O
  • Start with the element that is farther left on the periodic table (Hydrogen).
  • The rules for the second element still apply, suffix of “ide”.
  • Different is that the elements require prefixes depending on how many are in the compound.
  • So water’s chemical name is dihydrogen monoxide.
  • Prefix mono is only used for the second element.
  • “a” or “o” is left off of the prefix when used with an element starting with a vowel
diatomic molecules
Diatomic Molecules
  • Atoms can share electrons with the same atom.
  • These molecules have two of the same atoms joined by a covalent bond.
  • Since there are two of the same atoms the word diatomic is used. (“di” meaning two)
  • Seven elements exist as diatomics:
    • Hydrogen
    • Oxygen
    • Nitrogen
    • Fluorine
    • Chlorine
    • Bromine
    • Iodine
ionic compounds1
Ionic Compounds
  • Ionic compounds form large structures called lattices
  • Attraction between oppositely charged ions is strong.
ionic properties
Ionic Properties
  • Characteristics of an ionic compound:
    • Tend to have relatively high melting and boiling points because of the large amount of energy is needed to break the strong force of attraction in an ionic bond.
    • Conduct electricity when they are liquid or when they are dissolved in water. Melting or dissolving allow ions to move freely. In a solid state the ions are not able to move and therefore cannot conduct electricity.
molecular compounds1
Molecular Compounds
  • Bonds within the molecule are strong but forces of attraction between the molecules is weak.
molecular properties
Molecular Properties
  • Characteristics of a molecular compound:
    • Have relatively low melting points because little energy is needed to break the forces of attraction between molecules.
    • Relatively soft
    • Tend not to conduct electricity when they are in solid or liquid state. Do not conduct when dissolved in water because ions are not formed.
  • An electrolyte is a substance the dissolves in water to produce a solution that conducts electricity.
  • Ionic substances tend to be electrolytes.
  • Molecular substances tend to be non-electrolytes.
what is a reaction
What is a Reaction
  • Two atoms can be mixed together to create a new compound.
  • Can also mix two compounds to create new compounds.
  • New substances are produced in each case with properties different from the original substances.
dalton s atomic theory
Dalton’s Atomic Theory
  • All matter is made up of small particles called atoms
  • Atoms cannot be created, destroyed, or divided into smaller particles
  • All atoms of the same element are identical in mass and size, but different in mass and size from atoms of other elements
conservation of mass
Conservation of Mass
  • Law of conservation of mass states:
    • The total mass of the reacting substances (the reactants) is always equal to the total mass of the resulting substances (the products).
  • This means that the total number of atoms of each element must remain the same throughout the reaction
writing chemical equations
Writing Chemical Equations
  • Lets look at how water is formed.
  • First we can look at the word equation.

Hydrogen + Oxygen  Water

  • Reactants on the left side of the arrow
  • Products on the right side of the arrow
  • Plus sign means “reacts with”
  • Arrow means “produce”
skeleton equation
Skeleton Equation
  • For water the skeleton equation is:

H2 + O2 H2O

  • According to Conservation of Mass the reactants and products must have the same number of atoms.
  • Therefore we must write a balanced chemical equation.
  • This is done by adding coefficients in front of the compounds.
balancing the reaction
Balancing the Reaction
  • Important: Subscripts CANNOT be changed!!!!!!
  • Hydrogen is balanced so you must balance the oxygen by adding another water molecule

H2 + O2 2H2O

  • Now hydrogen is not balanced so we must add another hydrogen molecule

2H2 + O2 2H2O

check the atoms
Check The Atoms
  • Add up all the atoms on both reactants and products to make sure the amounts are the same.
what about state
What About State
  • The state of the chemicals is important to the chemical reaction.
  • State can be added by putting an abbreviation in parentheses after each chemical.
final balanced equation
Final Balanced Equation
  • Hydrogen and Oxygen are both gases
  • Water is a liquid
  • Final balanced equation with the states of all components is:

2H2(g) + O2(g)  2H2O(l)

  • You will not be required to give the states for an equation
rules for balancing
Rules for Balancing
  • Write the skeleton equation.
  • Pick an element within a molecule to balance first.
  • Treat polyatomic ions as a whole and not as individual atoms.
  • Balance more complicated molecules first.
  • Balance the individual elements last.
  • Leave Oxygen and Hydrogen until the end.
  • You cannot change the subscripts.
  • Now lets try balancing a few more chemical reactions:

Sodium + Water  Sodium Hydroxide + Hydrogen Gas

Copper + Silver Nitrate  Copper(II) Nitrate + Silver

Calcium Nitrate + Sodium Hydroxide  Calcium Hydroxide + Sodium Nitrate

synthesis reactions
Synthesis Reactions
  • Two or more reactants combine to produce a new product.

X + Y  XY

  • Example is water

2H2 + O2 2H2O

  • Others are the formation of table salt

2Na + Cl2 2NaCl

  • or the rusting of iron

4Fe + 3O2  2Fe2O3

decomposition reactions
Decomposition Reactions
  • A compound breaks down into two or more compounds or elements.

XY  X + Y

  • Example is water

2H2O  2H2 + O2

  • Another is the decomposition of table salt

2NaCl  2Na + Cl2

both put together
Both put Together
  • Common reaction that can demonstrate both synthesis and decomposition is pop.
  • Pop is carbonated water (Carbonic Acid)

Synthesis: CO2 + H2O  H2CO3

  • This causes the pop to fizz and bubble
  • When pop goes flat or is shaken up the opposite happens.

Decomposition: H2CO3  CO2 + H2O

  • The CO2 is released as a gas and water is left over.
single displacement
Single displacement
  • One element of a compound is replaced by a new element to form a new compound.

X + YZ  YX + Z

  • Example:
  • Sodium reacting with calcium chloride

2Na + CaCl2  2NaCl + Ca

  • The sodium takes the place of the calcium in the calcium chloride compound and makes sodium chloride.
double displacement
Double displacement
  • The elements of two different compounds exchange places, forming two new compounds.

WX + YZ  YX + WZ

  • Example:
  • AgNO3+NaCl  AgCl + NaNO3
  • A combustion reaction is when oxygen combines with another compound to form water and carbon dioxide.
  • These reactions are exothermic, meaning they produce heat.
  • The compound reacting with oxygen must have a carbon.
  • Example: Octane (component of gasoline)
  • __C8H18 + __O2 __CO2 + __H2O
  • 2C8H18 + 25O2 16CO2 + 18H2O
carbon compounds
Carbon Compounds
  • Organic chemistry is the study of compounds that contain carbon.
  • The study of all other compounds is known as inorganic chemistry.
  • Some carbon compounds are not considered organic (carbon dioxide, carbon monoxide and ionic carbonates).
  • Hydrocarbons are organic compounds that only contain carbon and hydrogen atoms.
  • The main source of hydrocarbon are crude oil and natural gas.
  • The combustion reaction of hydrocarbons is exothermic – lots of energy is released.
  • Thermal energy is given off that can warm our homes and provide energy for transportation.
complete combustion
Complete Combustion
  • Can have different levels of combustion depending on how much oxygen is present during the reaction.
  • When the hydrocarbon is burned in a plentiful supply of oxygen a complete combustion occurs.

Hydrocarbon + oxygen gas → carbon dioxide + water

incomplete combustion
Incomplete Combustion
  • When the hydrocarbon is burned in a poor supply of oxygen an incomplete combustion – not as much heat generated.
  • In an incomplete reaction there is carbon dioxide and water but also carbon (soot) and carbon monoxide.

Hydrocarbon + oxygen gas → carbon dioxide + water + carbon + carbon monoxide

  • Can be very dangerous!!!!
  • Synthesis
  • Two or more compounds come together and produce one compound.
  • Decomposition
  • One compound breaks down into two or more compounds.
  • Single Displacement
  • One element replaces another in a compound to produce a new compound.
  • Double Displacement
  • Two elements switch places with each other to produce to new compounds
  • Combustion
  • Oxygen combines with another compound to form water and carbon dioxide.
determination of acids and bases
Determination of Acids and Bases
  • Most solutions of acids and bases are clear and colourless.
  • To determine whether a solution is an acid or base an indicator is used.
    • Chemical that changes colour as the concentration H+ (aq) or OH- (aq) ions change
    • Two common indicators are phenolphthalein and litmus.
    • Phenolphthalein clear in the presence of an acid turns dark pink in the presence of OH+ ions
  • Red litmus paper turns blue for bases
  • Blue litmus paper turns red for acids
  • Red Acid and Blue Basic
ph scale
pH Scale
  • Litmus cannot determine how strong the solution is.
  • The pH scale measures the acidity of the solution. In other words it can be used to determine the concentration of H+ ions present in the solution.
  • Scale is from 0 to 14 with 7 being a neutral solution
  • 0-7 acidic solution and 7-14 basic solution
  • Water contains H+ and OH- ions but only a few water molecules ionize
    • H2O(l) ↔ H+(aq) + OH-(aq)
    • Neutral water contains an equal number of ions.
    • pH of water is 7
  • Two factors determine how many H+ ions are in the solution.
    • Concentration: Amount of pure acid that is dissolved per one litre of water.
    • Ionization: The process of molecules becoming ions.
  • Percent Ionization is the number of molecules that will ionize for every 100 molecules that dissolve.
  • Complete Ionization – Strong Acids and Bases.
  • Partial Ionization – Weak Acids and Bases.
  • Strong acids and bases are very hazardous.
  • Reaction between an acid and a base.
  • The resulting products are a salt and water.
  • Acid + Base  Salt + Water
  • A salt is an ionic compound
  • The combination of different acids and bases produces different salts.
  • Table salt, NaCl, can be produced from the following reaction.
  • HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
swimming pool
Swimming Pool
  • The process of neutralization is used to maintain pools.
  • Chlorine is used in pools to kill bacteria and algae.
  • Since chlorine gas is dangerous(corrosive and toxic), hypochlorous acid (HOCl) is used to keep pools clean.
  • To produce hypochlorous acid you must bubble chlorine gas through water.
  • A salt of HOCl can be produced by reacting with NaOH
  • Ideally pH should be between 7.2 and 7.8. If the water becomes too acidic (low pH) than a base (sodium carbonate) is used to neutralize the water.
  • If the pH becomes too high than hydrochloric acid can be used to remove the excess hydroxide ions.
acid rain
Acid Rain
  • When carbon dioxide go into the atmosphere it reacts with rainwater to produce carbonic acid.
  • CO2(g) + H2O(l)  H2CO3(aq)
  • This is how acid rain is produced
  • Can you think of a reaction we looked at before that involved carbonic acid.
  • Carbonated beverages contain carbonic acid.
  • So pop can’t be good for you right.
  • Uncontaminated rainwater has a pH of about 5.5
  • Rainwater with a pH below 5.5 is considered to be acid rain