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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

Matter

Mixture

Pure Substance

Homogeneous (solution)

Heterogeneous (mechanical mixture)

Element

Compound

classification of matter alternate
Classification of Matter (Alternate)

Matter

Homogeneous

Heterogeneous

Mixture

Pure Substance

Element

Compound

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

A

X

Z

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

23

Na

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

Metals

Metalloids

Non-Metals

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
Isotopes
  • 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:

H

He

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.

C

N

B

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.

Na

Cl

[ ]+

[ ]-

Na

Cl

example1
Example
  • 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)
examples
Examples
  • 1. Iron(III) Oxide (Rust)

Fe3+ O2-

Fe3+ O2-

O2-

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

Cu Cl-

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.

O

H

H

O

H

H

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.
prefixes
Prefixes
  • 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.
electrolytes
Electrolytes
  • 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.
examples1
Examples
  • 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
combustion
Combustion
  • 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
Hydrocarbons
  • 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!!!!
summary
Summary
  • 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
properties
Properties
  • 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.
neutralization
Neutralization
  • 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