Unit 2: Matter and Chemical Change

Lesson 1: Properties of Matter

All chemicals should be handled very carefully. WHMIS (Workplace Hazardous Materials Information System) has been developed to provide guidelines for handling, storage and disposal of reactive materials (chemicals).

Some materials are CAUSTIC - will burn, corrode or destroy organic tissue

MATTER: Matter is anything that has mass. All matter is described in terms of its properties

PROPERTIES - are characteristics you can use to describe or identify different substances. Ex. Color, luster, state, taste, melting point, and behavior

Properties can be classified as:

1. QUALITATIVE PROPERTIES- Those properties, which describe a quality of matter that has no numerical value (no number)- Usually involves one of the sensesEx. Taste, odor, texture, color, luster

2. QUANTITATIVE PROPERTIES -Those properties, which describe a quantity of matter- Have a number associated with the propertyEx. Melting point of water is 00 C Freezing point, boiling point, number of legs, density etc.

CHEMICAL PROPERTIES – those properties that involve the formation of a new substance. Chemical properties cannot be tested without destroying the substance.

Ex. Rust is formed when iron reacts with oxygen

Magnesium burns to produce a white powderPaper burns to produce carbon dioxide and water

PHYSICAL PROPERTIES – do not involve the formation of a new substance. Physical properties can be tested without destroying the substance.

Ex. Melting point – temperature at which something changes from a solid to a liquidBoiling point – temperature at which something changes from a liquid to a gasMalleability – can be pounded or rolled into sheetsDuctility – can be stretched into a wireDensity – the amount of mass in a given volume of a substanceConductivity – ability of a substance to conduct electricitySolubility – ability to be dissolved in another

The Particle Model of Matter states that:

1. All matter is made up of extremely tiny particles.2. Each pure substance has its own kind of particle, different from the particles of other pure substances.3. Particles attract each other. 4. Particles are always moving.5. Particles at a higher temperature move faster on average than particles at a lower temperature.

All matter can be classified according to its state as a

1. Solid – has a definite shape and a definite volume- Particles are close together and the forces between the particles are strongest

2. Liquid - has a definite volume but takes the shape of its container- Are said to be fluid because they flow- Particles are further apart and the forces between the particles are weaker

3. Gas - has neither a definite shape nor a definite volume- Expands to fill the container- Gases are fluids because they also flow- Particles are furthest apart and the forces between the particles are very weak

CHANGES OF STATE

1. Melting 2. Freezing3. Boiling/Vaporizing4. Condensation

5. Sublimation

Classification of Matter

Matter can also be classified according to its composition:

Matter

Pure Substances Mixtures

Elements Compounds Mechanical Solutions

Mixtures (homogeneous) (heterogeneous)

Suspension Colloids

PURE SUBSTANCES have properties that are always the same.

Ex. Table salt is a white solid that melts

at 801 °C and boils at 1465 °C.

Water, vinegar, sugar etc.

Elements

-Pure substances that cannot be broken down into any simpler substances - most elements are solids but several are gases and two are liquids-All of the elements have been arranged on the Periodic Table according to certain properties. - Contains only one type of atom Ex. Silver, Oxygen, Iron, Carbon,

Mercury

Compounds- Pure substances that contain two or

more different elements, combined in a

definite fixed proportion.

- Can be broken down chemically into different substances since it is made up of different kinds of atoms

Ex. Water - 2 hydrogen + 1 oxygen ---> H2O

Salt - 1 sodium + 1 chlorine ---> NaCl

MIXTURES contain at least two different substances. Properties are variable

Solutions -Homogeneous are a mixture in which one substance is dissolved in another. It is a uniform mixture that appears the same throughout.

Ex. Salt water, apple juice, and air

Mechanical - Heterogeneous are mixtures that do not appear the same throughout. The different components are visible.

Ex. Soil, chocolate chip cookies, chicken noodle soup

Suspensions – Heterogeneous mixture made of large particles that are uniformly mixed but will settle if left undisturbedEx. Sand in water, powdered chalk

in waterColloids - Heterogeneous mixture composed of fine particles evenly distributed throughout a second substance Ex. Hair gel Emulsions - Type of colloid in which liquids are dispersed in liquids

- Many will separate quickly to form layers of the original liquids (oil and vinegar)

Lesson 2: Chemical and Physical Changes

Physical Changes - No new substance produced- Change of size, shape or state - Ex. Cutting, freezing, molding, boiling, dissolving

Chemical Changes- Starting material is used up - New substance formed with different properties - Atoms are rearranged to form new molecules- Changes cannot be reversed - Ex. Cooking an egg, rusting, burning

Evidence of a Chemical Change

(a) Color Change

(b) Gas Formed – bubbling

(c) Solid material, called a precipitate is formed.

- two solutions are combined and a solid is formed

(d)Energy Change - energy is the ability to do work

- Ex. Light, heat, mechanical, sound, electrical

There are 2 types of energy change:1. Endothermic – energy is required.

- Energy is added to the starting materials

- Ex. cooking

2. Exothermic – energy is released.- Ex. Burning

Identify the following substances as pure substances (element / compound) or as a mixture (homogeneous / heterogeneous):

1. zinc 6. vinegar2. carbon dioxide 7. tossed

salad3. orange juice 8.

aluminum4. nitrogen 9. kool-aid

5. sugar 10. windex

Lesson 3: History of the Atomic Theory

Aristotle: (350 BC) – Greek philosopher

- Believed that everything was made of 1. Earth (dry and cold)2. Air (wet and hot)3. Fire (dry and hot)4. Water (wet and cold)

Robert Boyle: (1660’s) – England

Recognized that elements could be combined to form compounds

Lavoisier : (1770-1780) – France

1. Defined elements as pure substances that cannot be decomposed (broken down into simpler substances)

2. Developed a system for naming chemicals, so that all scientists could use the same words

3. Identified 23 pure substances as elements4. Discovered that in a chemical change, “the

mass of the new substances is always the same as the mass of the original substances” – LAW OF CONSERVATION OF MASS.+

John Dalton : (1808) – England

Atomic Theory:

1. All matter is made up of small particles called atoms2. Atoms cannot be created, destroyed or divided into smaller particles3. All atoms of the same element are identical in mass and size. Atoms of one element are different in mass and size from the atoms of other elements4. Compounds are created when atoms of different elements link together in definite proportions.

Dalton’s Theory led to the current definitions:

Element – a pure substance made up of one type of particle, or atom. Compounds – pure substances that are made up of 2 or more elements chemically combined together. Compounds can be broken down into elements again by chemical means.

J.J. Thompson: (1897) England

- Raisin bun model (plum pudding)

- Atom is a sphere, which is positive, with negative electrons embedded in it like raisins in a bun

Ernest Rutherford: (1911) McGill University, Canada

- Atoms have a nucleus which is positive - Most of the atom is empty space occupied by the moving negatively charged electrons- Proposed the existence of protons in a

nucleus

Neils Bohr: (1913) – Danish

-Electrons move in circular orbits around the nucleus- Like a miniature solar system

James Chadwick: (1932) - showed that the nucleus must contain heavy

neutralparticles to account for all of the atom's mass

- proposed the existence of neutrons

Lesson 4: Element Symbols- All elements have been given an atomic symbol

(a) A single capital letter – O – oxygen(b) Capital letter & a lower case letter – Co – cobalt(c) Capital letter & 2 lower case letters – Uun – ununnilium

- In the 1860’s Dmitri Mendeleev, a Russian chemist arranged the elements in order of increasing ATOMIC MASS and created the PERIODIC TABLE

- ATOMIC MASS is the average mass of an atom of an element

Ex. Oxygen = 16.00 g/mol

- Mendeleev found that the properties of the elements repeated at definite, or periodic intervals (ex. Lithium, sodium and potassium have similar properties so he placed them in the same family or vertical row)

- He left blanks in the table where he predicted elements should be and predicted what their properties would be, based on where they were on his table

- After the development of atomic theory, the periodic table was rearranged in order of increasing ATOMIC NUMBER

- ATOMIC NUMBER is the number of protons an element has in its nucleus

Ex. Fluorine – atomic number = 9, therefore it must have 9 protons in its nucleus.

The Periodic Table contains a lot of information about the different atoms. For example:

Atomic number Symbol

Name

Atomic Mass

4 Be

Beryllium

The horizontal rows on the periodic table are called Periods.

The vertical rows on the table are called Groups or Families.Elements in the same family have similar properties (behave in a similar manner)There are 18 Groups or Families.

The key on the periodic table will indicate the state of each element. Ex. White box – solid

Grey box – gasBlack box - liquid

- All elements can be classified as metals, non-metals or metalloids depending on their properties

Metals Non-metals- Found to the left of the - located to the right of staircase line of the staircase line- 80% of all elements - 20 % of all elements- Lustrous (shiny) - dull- Ductile (stretched into wire) - non-ductile- Malleable (hammered/shaped)- brittle- Conduct electricity - non-conductors- All solids, except mercury - Mostly gases, some metals, 1 liq.- Ex. Sodium, iron Ex. Oxygen, bromine

Metalloids - these elements have properties of both metals and non-metals

Ex. Silicon – shiny like a metal, poor conductor like a non-metal

There are 4 special named groups in the table:

Group 1 – Alkali Metals- Most reactive metals- Never found in pure form in nature

Ex. Lithium, sodium, potassium

Group 2 – Alkaline Earth Metals- React fairly vigorously with some substances

Ex. Magnesium, calcium, barium

Group 17 – Halogens- Most reactive non-metals

Ex. Fluorine, Chlorine

Group 18 – Noble Gases- Most non-reactive elements- Used to be called “Inert” gases until 1963

when a

Canadian chemist at UBC, made some of them react- Different noble gases produce different colors

Ex. Argon – blueHelium – yellow-white

Lesson 5: ATOMIC STRUCTURE

Atom - the smallest part of an element (which retains the chemical and physical properties of the element). Atoms are made up of 3 sub-atomic particles

1. Electron (e)

-Smallest particle in an atom-Has a negative charge

-Located in the extra nuclear region of the atom - outside the nucleus

2. Proton (p)

-Has a large mass -Has a positive charge

-Located inside the nucleus

3. Neutron (n)

-Same mass as a proton-Has a neutral charge (no charge)-Located inside the nucleus

Nuclear Notation

- Atomic number is the number of protons in the nucleus

- The number of protons equals the number of electrons in a neutral atom (#p = #e)

- Atomic Mass Number is the total number of protons and neutrons in the nucleus

And,

Atomic # = #p = #e

Example:

Find the number of protons, electrons and neutrons for iron and sodium.

Fe Atomic # = 26Atomic mass = 55.85 = 56 (round the

mass)Therefore: # of p = 26

# of e = 26# of n = 56 – 26 = 30

Note: when finding the number of neutrons we round the atomic mass to the nearest whole number.

Na Atomic # = 11Mass # = 22.99 = 23

p = 11e = 11n = 23 – 11 = 12

Number of neutrons = mass # - atomic #

Au Atomic # = 79Atomic mass = 196.96 = 197

p = 79e = 79n = 197 – 79 = 118

Lesson 6: Bohr’s Model of the Atom

- Bohr’s model states that electrons can be found only in certain energy levels or orbits around the nucleus

- He also stated that only a certain maximum number of electrons are allowed in each orbit.

Orbit # Max. # of electrons

1st 2 2nd 83rd 84th 185th 186th 32

When one orbit is filled the remaining electrons go to the next orbit – you cannot exceed the maximum allowed.

We can draw the Bohr diagram for any element. It musthave a nucleus showing the number of protons and neutrons and circles outside the nucleus showing the number of electrons.

Reminder: # of protons = # of electrons = atomic #

e.g. Draw the Bohr model for the following elements:

a) Lithium

Step 1 – Look up the atomic numberIt’s 3.

So, # of p = 3# of e = 3

Step 2 – Look up the atomic mass.

It’s 6.94 = 7 (round to the nearest whole #)

Find the number of neutrons.

Reminder: # of n = atomic mass – atomic #

So, # of n = 7 – 3 = 4

Step 3 – Draw the diagram.1st orbit

#p=3#e=3#n=4Atomic # =3Atomic Mass=7

The Bohr model diagram can be simplified – we can use lines instead of circles – this is called Electron Energy Level Representation.

Electron Energy Level Representations (EELR)Ex. Zinc - Atomic # = 30

Mass # = 65

Therefore, #p = 30 #e = 30 #n = 65 – 30 = 35

128 These must add

up to 30.82e

P=3n=4

P=30 n =35

Lesson 7: Molecules and Compounds

Molecule

– a particle formed when two or more atoms combine

* can be 2 atoms which are the same ex. H2, or O2

OR * can be 2 or more different atoms ex. CuS, NaCl, CO

Compound is a pure substance made of 2 or more different elements- Compounds can be broken down into simpler substances- ELECTROLYSIS – use of electricity to separate a chemical compound into its elements

Ex. Water –broken down into hydrogen and oxygen

- CHEMICAL BONDS hold elements together

- If elements SHARE electrons to form a bond it is called a MOLECULAR BOND --

MOLECULAR ELEMENT or MOLECULAR COMPOUND – Non-metals share electrons to form molecular compounds.

E.g. CO2

If atoms transfer electrons from one atom to another to form a bond it is called an IONIC BOND -- IONIC COMPOUND – metals transfer electrons to non-metals to form ionic compounds.

E.g. NaCl - Sodium chloride

- CHEMICAL FORMULAS use symbols and numbers- If only one atom is represented, no numbers are used- if there is more than one of that type of atom present a small number written below the line is used to tell us the number of that type of atom. This is called a SUBSCRIPT.

E.g. H2O - One water molecule is made up of

2 atoms of Hydrogen and 1 Oxygen.

NaCl - sodium chloride 1 – sodium 1 – chlorine

H2 - hydrogen 2 – hydrogen

C12H22O11 - sucrose 12- carbon 22 – hydrogen11 – oxygen

Cu(NO3)2 - copper nitrate 1 – copper2 – nitrogen 6 – oxygen

All pure substances can be identified in two ways: * Element or compound* Atom or molecule

NaCl - compound/moleculeH2 - element/moleculeC6H12O6 - compound/moleculeZn - element/atomCu(NO3)2 - compounds/molecule

Lesson 8: Molecular Elements and Compounds

Diatomic Elements – Molecules made of 2 atoms of the same element- All of the halogens, plus oxygen, hydrogen and nitrogen are diatomic elements

H2, O2, F2 Cl2, N2, Br2

Molecular Compounds- Formed when 2 non-metals share electrons- Most molecular compounds have low melting and boiling points; therefore they are found as solids, liquids and gases at room temperature- They are poor conductors

Naming Binary Molecular Compounds

IUPAC – International Union of Pure and Applied

Chemistry – Determines how compounds are named - followed by chemists around the world

Step 1. Write the entire name of the first elementStep 2. Change the ending on the name of the second element to – ide Step 3. Use a prefix to indicate the number of

each type of atom in the formula.

Prefixes are:1 – mono (only used for the second

element)CO – carbon monoxide

2 – di 3 – tri4 – tetra5 – penta6 – hexa7 – hepta

8 – octa9 – nona10 – deca

Ex. P2O5 – diphosphorous pentaoxide SiO - silicon monoxide SCl2 - sulphur dichloride

NO2 – Nitrogen dioxide

N2O – Dinitrogen monoxide

N2O4 – Dinitrogen tetraoxide

Writing Molecular Formulas

1. Write the symbol for the elements in the same order as they appear in the name.2. Use subscripts to indicate the numbers of each type of atom.

Ex. Carbon tetrabromide - CBr4 Triarsenic hexasulphide - As3S6

We use small symbols in parentheses after the formula for each compound to indicate the state of matter

(s) - solid - NaCl(s)(l) - liquid - H2O(l)(g) - gas - CO2 (g)

Prefixes:

1 = mono 6 = hexa2 = di 7 = hepta3 = tri 8 = octa MEMORIZE4 = tetra 9 = nona5 = penta 10 = deca

Assignment 11: Name_________________________

Name or give the formula:

1. Silicon dioxide ____________________________________2. Sulphur monoxide_________________________________3. OF2 _________________________________________4. SiBr4 ______________________________________5. PH3 _______________________________________6. N2O _______________________________________

7. CO ________________________________________8. NBr3 _______________________________________9. P2I3 _________________________________________10. SO3 _________________________________________11. N2O4 _____________________________________________12. Tetraphosphorous hexaoxide _________________________________13. Dinitrogen tetraoxide _______________________________________14. Heptasilicon monobromide _____________________________________15. Octaboron decaiodide _________________________________________

16. B2O3 ______________________________________________17. BrF7 ______________________________________________18. N3O6 _____________________________________________19. H2Cl5 _____________________________________________20. Triselenium diastatide _________________________________21. Diarsenic pentaoxide _____________________________________22. Sulphur trioxide _________________________________________23. C3O2 ___________________________________________________24. C2H6 ___________________________________________________25. As3Br7 __________________________________________________26. SO2 _____________________________________________________

27. Selenium monoxide __________________________________________28. Diboron trioxide _____________________________________________29. PF3 ________________________________________________________30. P2O5 __________________________________________

__________________________________________________31. P4O10 ________________________________________32. Arsenic trifluoride ______________________________33. BrF7 __________________________________________

Lesson 9: Ionic Compounds- Atoms that gain or lose electrons to become stable are called IONS.- If they gain an electron they have more negative charges than positive charges so they have a slight negative charge- Non-metals gain electrons.

- If they lose an electron they have less negative charges than positive charges so they have a slight positive charge- Metals lose electrons.

Ionic Compounds

- Made up of a metal bonded to a non-metal – electrons transfer from the metal to the non-metal.- All solids at room temperature- Separate into positive and negative ions when they dissolve in water- The ions conduct electricity

Naming Ionic Compounds

1. The name includes both elements in the compound, with

the name of the metallic element first.

2. The non-metallic element is second. Its ending is

changed to – ide.3. No prefixes used in naming.

Ex. CaCl2 - calcium chloride (1 calcium/2 chlorine)

Na2S - sodium sulphide(2 sodium/1 sulphur)

Fe2O3 - iron(iii) oxide (rust)

Lesson 10: Chemical Reactions

Chemical Reactions: formation of a new substanceChemical bonds are broken and new bonds formed

Reactants Products Starting materials become End materials

Reactants: Any substance that is used up in the reaction Products: Any substance that is produced in the reaction

Word equation: -Gives the names of all the reactants (separated by a + sign)-Arrow points to the names of all the products (separ. by + sign)

Write Word Equations for the following reactions:

(a) When sodium reacts with chlorine, sodium chloride (salt) is produced.

Sodium + Chlorine Sodium chloride

(b) Hydrogen gas and zinc chloride are produced when a piece of zinc metal is dropped into a beaker of hydrochloric acid.

Zinc + Hydrochloric acid Hydrogen + Zinc chloride

(c) Potassium iodide is decomposed to produce potassium metal and iodine.

Potassium iodide Potassium + Iodine

Law of Conservation of Mass:In a chemical reaction the total mass of the reactants is always equal to the total mass of the products.

In a chemical reaction mass is neither gained nor lost. Molecules may be broken apart and new ones may be formed, but the atoms in the products are the same ones that were in the reactants

Ex. Start with 100 grams of reactants - end up with 100 grams of product.

Start out with 20 Hydrogen and 10 oxygen - end up with 20 hydrogen and 10 oxygen

Using numbers called COEFFICIENTS in front of the elements and compounds in the reaction balances chemical reactions.

2 H2O --- 2 H2 + 1 O2

Reaction Rate – A measure of how fast a reaction occurs- Some reactions are naturally fast others are slow

We can influence the rate in several ways:

1. Temperature – The higher the temperature the faster the

rate- Molecules move faster at higher

temperatures- Molecules collide more often and form new substances more quickly- The faster the rate, the less time needed for the reaction

2. Concentration- Refers to the amount of solute present in a specific amount of solution - the higher the concentration, the faster the reaction

3. Surface Area - For 2 substances to react, they must come into close contact - the greater the surface area, the more contact the two substances have therefore the faster the reaction- Can increase surface area by grinding up a chemical

4. Stirring - Increases the chances of collisions,

therefore speeds up the reaction rate

Lesson 11: Catalysts and Inhibitors

Catalyst- A substance that speeds up the rate of a reaction (without being changed itself).

Ex. saliva – acts as a catalyst to break down starch

Inhibitors- Substances that slow down chemical

reactions

Ex. Added to some foods and medicines to slow down their decomposition

Corrosion

Corrosion is a chemical reaction. It is the “eating away” of a metal as it reacts with other substances in the environment.

Corrosion of iron is called rusting

4 Fe(s) + 3 O2(g) 2 Fe2O3(s)iron + oxygen produces rust

Different metals corrode at different rates. Iron corrodes quickly. Gold does not corrode at all. Aluminum and copper corrode only on the surface. Corroded materials lose their strength. Rusting is sped up by high temperature, surface area, and the presence of water, air, salt or acid.

If the metal is totally protected from air or water, rusting cannot occur.

Rust Protection

1. Paint.

2. Rust Check – spray with an oil film to keep air and water away

3. Galvanization - coat it with zinc - Zinc is more resistant to corrosionEx. Galvanized nails

4. Electroplating – covering a metal with another metal by using electrolysis

-Ex. Bumpers are coated with a thin layer of chromium - chromium improves the hardness, stability, and appearance.

Combustion (Burning)

- Chemical reaction that occurs when oxygen reacts with a substance to form a new substance- Combustion is exothermic- Oxygen is always one of the reactants – no oxygen no combustion (no fire).

- Carbon dioxide and water vapor are the products of combustion when one of the reactants contains “carbon”

Ex. Methane - CH4

CH4 + 2 O2 CO2 + 2 H2O

Methane + oxygen Carbon dioxide + water

Identification Tests for the Products of Combustion

1. Test for carbon dioxide

- Bubble the gas through limewater solution- If the limewater turns milky, the gas is carbon dioxide

2. Test of water- Touch the cobalt (II) chloride paper to the

liquid- If the paper turns pink, the liquid is water