Lesson 1 - Working With Chemicals

Safety is the number one concern WHMIS (Workplace Hazardous Materials

Information System) was established to standardize information and symbols about chemicals in our liveso WHMIS informs workers about the chemicals in

three ways1. Controlled products must have informative labels,

in both English and French on their containers

2. Each controlled product must have a MSDS (Materials Safety Data Sheet).

3. Workers who handle chemicals must complete an education program provided through their employer.

http://www.brocku.ca/oehs/safety/whmis_symbols.pdf

MSDS (Material Safety Data Sheet) was required to accompany every chemical bought and sold (p.7).

http://www.lindecanada.com/en/msds/linde/Argon__Liquid_EN.pdf

Classifying Matter

Matter is defined as anything that has mass and volumeMatter may be solid, liquid or gas.

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Matter

Mixture Pure Substances

Combinations of matter Matter that has a that can be separated definite composition.by physical means

Do not have a definite composition

Gravel Milk Gold Water (H2O)2

HeterogeneousMixtures(Mechanical mixtures)

Different components of the mixture are visible or layers can be seen.

Homogeneous Mixtures (Solutions)

Different components are not visible

Element

Cannot be chemically broken down into simpler substances

Compound

Two or more elements that are chemically combined

Can be separated chemically into simpler substances

Assignment:

1. State whether the following is a pure substance or a mixture:

a) sea water b) iron c) bronzed) 14k gold e) table salt f) oxygen

2. State whether the following mixtures are homogeneous or heterogeneous:

a) Oil-and-vinegar salad dressingb) Steelc) Cranberry juiced) Sugar dissolve in watere) Milkf) Antifreeze

3. State whether the following pure substances is an element or a compound:

a) Copper b) Water, H2Oc) Methane, CH4 d) Silver

4. Classify each of the following substances:a) Graphite, C b) Shampooc) Coffee d) Motor oil

5. Name the following elements:3

Au, S, Fe, Hg, W, Cu, At, K, Na, Pb, Zr, Mo, Ag, P,Ca, Cr, Ac, Ne, Fr, Sc, Ar, N, Mn, Be, Pt, Bi, Kr, CHf, Th, Cs, Po, U, He, Y, Ir, In, Rn, Ce, Pu, Sb, Os,

Assignment:

a) p. 11 # 1 – 4 (copy question first)

b) Find the number of protons, electrons, and neutrons for the following

elements:

Hf, Th, Pt, Po, Au, U, Bi, Cs, Pb, W

Oxidizing material – rusting caused by oxygen (ex. iron)

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Developing (History of) Atomic Theories

1.Dalton’s Atom (1766-1844)

Dalton’s Atomic Theory (Pg 12)- All matter is made up of small particles called

atoms- Atoms cannot be created, destroyed, or divided

into smaller particles- All atoms of an element are identical in mass

and size, but they are different in mass and size from the atoms of other elements

- Compounds are formed when atoms of different elements combine in fixed proportions

- Chemical reactions change the way atoms are grouped, but atoms themselves are not changed in reactions

- “billiard ball” model

http://www.rsc.org/chemsoc/timeline//pages/1803.html

2.J.J. Thomson (1856-1940)- English physicist- Atoms contain negatively charged electrons- Electrons are like raisins in a plum pudding or

“raisin bun” model.

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3.Ernest Rutherford (1871-1937)- Atom contains electrons and positively-charged

particles- Atom composed of

o A nucleus – a central region that is positively charged and contains most of the mass- protons are heavy positive particles within the nucleus

o Electrons – particles with a negative charge and are very light (compared to protons).- Electrons circle around the nucleus

o Empty space surrounding the nucleus is very large within which electrons move (planetary model).

o Rutherford also proposed existence of the neutron to account for the mass difference between hydrogen and helium

o Neutrons are heavy particles like protons but have no charge

o Isotopes are atoms of the same element that differ in mass (but are chemically alike). (element with different number of neutrons)

http://www.iun.edu/~cpanhd/C101webnotes/modern-atomic-theory/rutherford-model.html

4.Niels Bohr (1885-1962)

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- Electrons exist only in certain energy levels or orbits around the nucleus

- Only a certain number of electrons can exist in each energy level (orbit).

5.Modern Theory- Present day models of the atom are much more

complex- Electron energy levels are divided into

sublevels.- Neutrons and protons are made of even smaller

particles called quarks.

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- or e)

-smallest particle in an atom-has a negative charge-located in the extra nuclear region of the atom

2. Proton (e+ or p)-Has a large mass

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-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

- The mass number also identifies the particular isotope.

Atomic # = #p = #e

Example:

Find the number of protons, electrons and neutrons 8

Number of neutrons = mass # - atomic #

for iron and sodium.

Fe Atomic # = 26Atomic mass = 55.85

p = 26e = 26n = 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

p = 11e = 11n = 23 – 11 = 12

Assignment:

Find the number of protons, electrons, and neutrons of the elements with atomic numbers 1 to 30 and 40-70.

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p. 11 #1-3 (copy questions or complete sentences)

Nuclear Notation Continued

1. One way to write isotopes of elements is:

12 6C where the top number is equal to the

atomic mass, and the bottom number is equal to the atomic number.

Atomic mass

12 6 C #p = 6

#e = 6Atomic number #n = 12 – 6 = 6

13 6C #n = 13 – 6 = 7

2. Another notation used is: e.g. Lithium–7 or Li - 7where 7 is equal to the atomic mass.

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From the table of elements we get the atomic number (which is 3).

Thus, #p = 3#e = 3#n = 7 – 3 = 4

p.23 #5 to 8p.24 # 1 to 4p. 37 (b, c) (copy question first or complete

sentences)

p. 38 - define the key terms (first column)

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Bohr’s Model of the Atom

According to Bohr’s model electrons exist only in certain energy levels or orbits around the nucleus

Only a certain number of electrons can exist in each energy level or orbit.

The 1st orbit can hold a max. of 2 electronsThe 2nd a max of 8

3rd 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:12

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.

#p=3#e=3#n=4

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P=3N=4

We can draw the orbits using a simplified version.

e.g. Cobalt ___ 9 ___ 8 ___ 8 ___ 2 e

Assignment: Draw the Bohr model of the atom for the elements: K, Cl, Mn, Zr, Ne, Ca, Sc, Ti, Fe, Al, Br, Ni, Cu, S, Ag, He, B, O, Na, Mg, Be, Ar, N, V,

1to10 and 15 to 30.

And p. 38 #1, 5, 6 (copy question or complete sentences)

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P = 27N = 32

Periodic Table

Periods are horizontal rows in which elements increase in atomic mass from left to right

Groups or families are vertical rows made up of elements with similar properties. There are 4 special named groups.

Group 1 – Alkali MetalsGroup 2 – Alkaline EarthsGroup 17 – HalogensGroup 18 – Noble gases or Inert gases

- ‘Staircase’ line separates metals from non-metals

- Metalloids border this line

Francium is the most reactive metal.Fluorine is the most reactive non-metal.

Valence Electrons

- The outermost occupied energy level (orbit) of an atom is called its valence energy level.

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- The electrons in the valence energy level (electrons in the last orbit) are

called valence electrons.

Electron dot diagrams or Lewis dot are useful ways to represent an atom.

In an electron dot diagram, the electrons in the last orbit are shown as dots placed around the symbol.

e.g.

Li

Bohr diagram

There is 1 valence electron (that is, 1 electron in the last orbit). •

So, the dot diagram will be Li e.g. Oxygen

#p = 8#e = 8#n = 8 ____ 6

____ 2e

16

#p = 3

#n = 4ee

p = 8n = 8

There are 6 valence electrons.

•• Electron dot or Lewis dot diagram • O ••

•

Assignment:

Draw electron dot diagrams for the following elements:K, Cl, Mn, Zr, Ne, Ca, Sc, Ti, Fe, Al, Br, Ni, Cu, S, Ag, He, B,

and p. 27 #9 to 12 p. 28 complete the table

Science Test Friday17

Assignment:

Draw electron dot diagrams for:

1. Scandium2. Fluorine3. Beryllium4. Vanadium5. Gallium

p. 38 – Define the key terms andp. 38 # 7, 8, 13

Quiz

Draw electron dot diagrams for:

1.Sc2.Na3.Chromium4.Ar

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Ions

- Ions are atoms, which have gained or lost electrons, in order to become more stable – it happens during chemical reactions.

- Ions always have a charge

- Positively charged ions have fewer electrons than protons – also called cations.

- Most metals form cations – that means they lose electrons

e.g. Li1+

Li loses an electron

Li1+

- Negatively charged ions have more electrons than protons – also called anions.

- Non-metals that form anions have a name ending in ‘ide’

e.g. chloride (Cl-), oxide (O-2 or O2-)

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3 6.941+

Li

Lithium

All non-metals gain electrons (that is, form anions).

Compounds

Compounds are formed when two or more elements are chemically combined.

- Noble gases with their 8 valence electrons are very stable elements – they usually don’t form compounds.

- Other atoms have different ways of becoming stable – they either gain or lose electrons when they form compounds.

- **Metals give up electrons to other atoms, forming cations.

- **Non-metals accept electrons, forming anions.

- **Non-metals may share electrons with other atoms.

e.g. Sulphur dioxide

non-metal non-metal

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Assignment p.36 #1-3, 5p.38 #6, 9, 15,16 (copy question or complete sentences)

Compounds

There are two basic types:1. Ionic2. Molecular

Ionic Compounds

- Ionic compounds formed from just two elements are called binary ionic compounds

- A metallic cation is joined to a non-metallic anion by an ionic bond.

- Ions of an ionic compound are arranged in a regular repeating pattern called crystal lattice.

Ionic compound – metal and non-metal joined chemically.

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In ionic compounds electrons are traded.

e.g. NaCl (p.30 – sketch Fig. 1.27 here)

Molecular Compounds

- Atoms which share electrons to become stable form molecular compounds (see p.32)

- These groups of atoms are called molecules- Atoms in molecules are joined by covalent

bonds.- All atoms in molecular compounds are non-

metals.

Molecular compounds – non-metal and non-metal joined chemically.

e.g. CO2 (p. 32 – sketch Fig. 1.29 here)

Assignment:

p.37 (i,j,k)p. 38 #13,15,16,17,18 (copy question)

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Investigation 1-A Pg 33

Check Your Understanding Pg 36

Read Pg 37

Review Pg 38

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Naming and Writing Binary Molecular Compounds

- When two (binary) non-metallic atoms join by a covalent bond we have a molecular compound.

e.g. Carbon dioxide

Rules for naming

1.The first element in the compound is the one most left on the periodic table.

2.The suffix ‘ide’ is attached to the name of the second element.

3. Prefixes are used to indicate how many atoms of each type are present in one molecule of the compound.

Prefixes:

1 = mono 6 = hexa2 = di 7 = hepta3 = tri 8 = octa MEMORIZE

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4 = tetra 9 = nona5 = penta 10 = deca

No “mono” is used with the first element.

e.g. Give the name or formula for each compound:

NO2 – Nitrogen dioxide

N2O – Dinitrogen monoxide

N2O4 – Dinitrogen tetraoxide

Nitrogen monoxide - NO

Dinitrogen pentaoxide – N2O5

Carbon dioxide – CO2

Assignment:

Name or give the formula:1. Silicon dioxide2. Sulphur monoxide3. OF2

4. SiBr4

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5. PH3

6. N2O7. CO8. NBr3

9. P2I3

10. SO3 11. N2O4

12. Tetraphosphorous hexaoxide13. Dinitrogen tetraoxide14. Heptasilicon monobromide15. Octaboron decaiodide16. B2O3

17. BrF7

18. N3O6

19. H2Cl5

20. Triselenium diastatide21. Diarsenic pentaoxide22. Sulphur trioxide23. C3O2

24. C2H6

25. As3Br7

26. SO2

27. Selenium monoxide28. Diboron trioxide29. PF3

30. P2O5

31. P4O10

32. Arsenic trifluoride26

33. BrF7

34. Hydrogen chloride35. N2O

And p. 44 #1- 4, p. 62 #1 (copy question first)

Binary Ionic Compounds

- Are composed of ions of one metal element and ions of one non-metal element joined by ionic bonds

Rules for naming

1. The first element in the name of the formula is the metal

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2. The second element, the non-metal, is named as an ion. The suffix ‘ide’must be present.

3. No prefixes are used.

e.g.

Fe2O3 – Iron oxideCuS – Copper sulfideKCl – Potassium chloride

p. 45 #5 - 7p. 46 #9, 10p. 47 #12p. 48 #3, 5 (copy question)

Writing Formulas for Binary Ionic Compounds

In an ionic compound the total number of positive charges must equal the total negative charges – the compound must be electrically neutral.

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This fact tells us how many of each atom is necessary to form a compound.

e.g. sodium chloride

Step 1 – use the table to find the charges on each ion (element)

Na1+ Cl1-

Step 2 – bring the two ions close together and see what the net charge is.

Na+Cl- the two charges are equal so the formula isNaCl.

Magnesium chloride

Mg2+Cl1-

Question: how many of each ion is needed so that the molecule is neutral.

Cl1-

Mg2+

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Cl1-

Therefore the formula is MgCl2

Chromium oxide Cr3+O2-

Cr3+O2-to balance the charges we use a

shortcut method – charges are “traded” across.

Cr2O3

Calcium oxide Ca2+O2- Ca2O2 CaO

Multivalent Cations (metals)

- Some atoms are able to form more then one cation. Ex. Ni2+ or Ni3+

- In the Stock system, the charge on the cation is written in brackets, as a Roman numeral after the name of the metal

Example

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Copper (II) oxide Cu2+O2- CuO

Tin (IV) fluoride Sn4+F1- SnF4

PbI2 Lead (II) iodide

Pb2+ I1-

Cr2S3 Chromium (III) sulfide

Cr3+S2-

Is this formula correct LiO Li1+O2-

No – correct formula is Li2O

p. 47 #11p. 48 #4, 5p. 49 #7, 9

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Polyatomic Ions

- Consist of two or more different atoms (covalently bonded) containing an overall charge. e.g. NO3

-

- Found in the box at the top of the table.

- All are negatively charged, except ammonium ion, and most names end in ‘ate’

- All act as non-metals except ammonium ion, NH4

+, which acts as a metal in compounds.- The name of the cation (metal) is followed by

the name of the anion (non-metal – negatively charged).

- When writing formulas, brackets must surround the polyatomic ion (when more than one is present – i.e. subscript is not 1).

Examples:

1. Potassium sulphate K1+(SO4)2- “trade” charges

K2(SO4) or K2SO4

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NH4NO3 Ammonium nitrate

Al(NO3)3 Aluminum nitrate

Sodium sulfate Na1+(SO4)2- Na2SO4

Na1+

SO42-

Na1+

Ammonium phosphate (NH4)1+(PO4)3-

(NH4)3PO4

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Gallium hydrogen carbonate

Ga3+(HCO3)1- Ga(HCO3)3

Assignment:

Practice Problems p. 52 #13-16Practice Problems p. 53 #17-18

p. 55 #1-3

Properties of Ionic Compounds

- In the solid state ionic compounds are crystalline- Ionic compounds have fairly high melting points- In the solid form they do not conduct electricity- In the aqueous (dissolved in water) form ionic

compounds are electrolytes – they conduct electricity

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Properties of Molecular Compounds

- Most molecular compounds have fairly low melting points – weak intermolecular bonds

- Non-electrolytes – do not conduct electricity- When dissolved in water most do not conduct

electricity (some do)

p. 55 #4, 5p. 80 #1 - 3, 6, 17,18, 20 (copy question)

Special Compounds and Elements

Special compounds – these compounds have special names, which do not follow the rules for naming.

Water H2O or HOHOzone O3

Ammonia NH3

Hydrogen Peroxide H2O2

Methanol CH3OHEthanol C2H5OH

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Sucrose C12H22O11

Glucose C6H12O6

Methane CH4

Diatomic and Polyatomic Elements

If these elements are FREE, that is ALL ALONE, they are written as:

H2 P4

N2 S8

O2

F2 MEMORIZECl2

Br2

I2

At2

For example, hydrogen has one electron and thus it wants to fill that orbit in order to become stable - so it will pair up with another hydrogen atom and they will share the two electrons – covalent bonding.

H • • H

Thus, hydrogen when it’s not in a compound but all-alone is written as H2.

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Assignment:p. 79 (a, b, c, e,)p. 80 – key termsp. 81 #25

Test tomorrow.

Properties of Acids and Bases

1. Acids- a substance that reacts and releases hydrogen

ions, H+(aq),in a water solution

- taste sour- form colourless solutions- conducts electricity- formula starts with Hydrogen

e.g. HCl - Hydrochloric acidH2SO4 - Sulphuric acid

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2. Bases- a substance that dissolves in water and releases

hydroxide ions, OH-

- bitter tasting- feel slippery- form colourless solutions- conducts electricity

e.g. NaOH - Sodium hydroxide

Indicators and pH

- an indicator is a chemical that changes a different colour in an acid vs a base

- litmus is red in acids and blue in bases- phenolphthalein is colourless in acids but pink in

bases- pH is a scale used to indicate the strength of the

acid or base- pH scale ranges from 0 – 14, - pH of 7 is neutral – pure water- pH less than 7 – acid- pH greater than 7 – base

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0 acid 7 base 14

p.80 #12,13p.81 #23,25 (copy questions)

Naming Acids

All acids start with hydrogen. Acids have special names, which derive from the following rules.

Chemical name Acid name

Hydrogen _______ide becomes Hydro______ic acid

e.g. HCl Hydrogen chloride Hydrochloric acid

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H2S Hydrogen sulfide Hydrosulfuric acid

Hydrogen _______ate _______ic acid

H2SO4 Hydrogen sulfate Sulfuric acid

HClO Hydrogen chlorate Chloric acid

Hydrogen ______ite _________ous acid

H2SO3 Hydrogen sulfite Sulfurous acid

HClO2 Hydrogen chlorite Chlorous acid

p. 70 #20 (a, b, c) #21 (a, b, c)#22,23

p. 71 #6p. 79 (e, k) and p. 135 #26-28

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Water

- the shape of the water molecule is

Oxygen end – slightly negative.

1050

Hydrogen end – slightly positive.

- has two covalent bonds but the electrons shared in these bonds are not shared equally

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- oxygen attracts the pairs of electrons closer to it- this creates an uneven distribution of charges or

partial charges

- the result is a polar molecule or dipole- the negative end or oxygen of one water attracts

the positive end or hydrogen of another – hydrogen bond

- hydrogen bonds are one kind of intermolecular force

- intermolecular forces are attractions between molecules

- intramolecular forces are attractions within molecules

Properties of Water

- the boiling point and melting points are higher in water than other similar substances – the need to break the hydrogen bonds

- it requires a great deal of energy to raise the temperature of water – strong intermolecular forces

- has a concave meniscus and shows capillary action – strong force of attraction between water and other molecules

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- ice floats in liquid water – due to the rearrangement of the hydrogen bonds in the solid creating a greater volume and lower density

- has a high surface tension – again due to the hydrogen bonds

Chemical Reactions

- chemical reactions occur when one or more substances change to form new substances

- also called a chemical change- the substances that change are called the

reactants- the substances formed are called products- evidence that a chemical change has occurred

could involve one or more of the followingo energy change – heat and/or light

exothermic – release energy endothermic – absorb energy

o odour changeo colour changeo formation of a gas – bubblingo formation of a solid – precipitate

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Predicting Solubility

- some ionic compounds are highly soluble in water while others have a very low solubility

- we use a solubility table to help determine whether a substance is soluble or not – back of table.

Step 1 – Locate one of the ions in the compound in the boxes across the top.

Step 2 – Look for the other ion in the two vertical boxes below.

If it is soluble write (aq) behind the compound to show that it is aqueous – it dissolves.

If it is slightly (low) soluble show that it does not dissolve by writing (s) behind the compound so that it is solid.

e.g. Determine if the following compounds are soluble or not by using the appropriate notation.

NaCl(aq)

Look for Na1+ or Cl1- across the top horizontal row.

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PbI2(s)

NH4OH(aq) CuSO4(aq)

p.90 #1, 2p. 93 #1- 4p.128 # 9

Law of Conservation of Energy

-Energy can be converted from one form to another, but the total energy of the universe remains constant (energy cannot be created nor destroyed).

- Breaking chemical bonds is an endothermic process (energy is used).

- Forming new chemical bonds is an exothermic process

1.When more energy is required to break bonds than is released when new bonds form, the reaction is endothermic

e.g. energy + water hydrogen + oxygen45

2.When less energy is required to break bonds than is released when new bonds form, the reaction is exothermic

Example: hydrogen + oxygen → water + energy

Lavoisier’s Law of Conservation of Mass

During a chemical reaction, the total mass of the reacting substances (reactants) is always equal to the total mass of the resulting substances (products).

Balanced Chemical Equations

- a balanced chemical equation shows that atoms are conserved in a chemical reaction(that is, the numbers of each atom must be equal on both sides of the equation).

- reactants are on the left side of the equation and products are on the right side

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- coefficients are used to balance a chemical equation (tells us how many molecules or atoms are needed in the reaction).

e.g.

H2 + O2 → H2O (this is called askeleton equation)

2H2 +1O2 → 2H2O (this is a balanced equation)

H + O → O + OH O H H H HHH

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http://funbasedlearning.com/chemistry/chemBalancer/

Assignment:

Model Problem 1 p. 99Model Problem 2 & 3 p. 100p. 101 #5p. 102 #5 (copy the EQUATIONS)

Types of Chemical Reactions

1. Formation Reactions or Simple Composition

- Two or more elements combine to form a new compound

Element + Element → Compound

X + Y → XY

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The reactions must be balanced.

e.g. Iron combines with oxygen to form iron (III) oxide.

4 Fe + 3 O2 → 2 Fe2O3

Copper reacts with chlorine to form copper (I) chloride.

2 Cu + 1Cl2 → 2 CuCl

p. 114 #3p. 115 #6 (Copy and balance)

2. Decomposition Reaction or Simple Decomposition

One compound breaks down into two or more elements

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compound → element + element + …

XY → X + Y

XYZ → X + Y + Ze.g.

2HCl → 1H2 + 1Cl2

2K2IO3 → 4K + 1I2 + 3O2

p. 127 (a-i,l,m) → complete sentences or copy the question.

p. 128 #2, 4, 9

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3. Single-Replacement Reactions

one element takes the place of another element in a compound- many involve the reaction between a metal and a

compound

element + compound → new element + new compound

A + BX → AX + B

AX + Y → AY + X

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Cu + 2AgNO3 → 2Ag + Cu(NO3)2

2NaBr + Cl2 → 2NaCl + Br2

Note: Metal replaces (switches with) a metal. Non-metal replaces a non-metal.

Mg + CuSO4 → Cu + MgSO4

When a metal reacts with water, the water formula is written as HOH (first H “acts” as a metal).

2Na + 2HOH → H2 + 2NaOH

OH1- (hydroxide ion)

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4. Double-Replacement Reactions

- Two different compounds react, forming two new compounds

compound + compound → new + new compound compound

AX + BY → BX + AY

Note: metal switches with a metal and a non-metal with a non-metal.

- a special kind of double-replacement reaction called neutralization is between an acid and a base

NaOH + HCl → NaCl + HOH

Base Acid

Ba(OH)2 + Na2SO4 → 2NaOH + BaSO4

53

p. 114 #2,3,7 p.134 #21 p.136 #40

5. Hydrocarbon Combustion

A hydrocarbon is an organic compound containing carbon and hydrogen (sometimes oxygen also)

When hydrocarbons are burnt in a plentiful supply of oxygen complete combustion occurs- The two products are always carbon dioxide and

water vapour

Hydrocarbon + oxygen → carbon dioxide + water

Hydrocarbon + O2 → CO2 + H2O

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When hydrocarbons are burnt in a poor supply of oxygen incomplete combustion occurs.

The products of this reaction are; carbon dioxide, water, carbon (soot) and carbon monoxide.

Carbon monoxide is an odourless, colourless and highly toxic gas.- CO binds 200x more strongly to hemoglobin in

the red blood cells than does O2

e.g.

CH4 + 2O2 → CO2 + 2H2O

2C2H6 + 7O2 → 4CO2 + 6H2O

p. 114 #1,4 p. 135 #27-29

Ammonia NH3

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Ethane C2H6

Glucose C6H12O6

The Mole

The mole is defined as the amount of substance that contains as many elementary entities (atoms, molecules, or formula units) as exactly 12 g of carbon-12, the most common isotope of carbon.

One mole of a substance has been determined to contain 6.02 x 1023 elementary entities of a substance (atoms, molecules). This number is called Avogadro’s number.

(Similar to dozen) dozen = 12mole = 6.02 x 1023

Atomic Molar Mass

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- is a weighted average of the mass of 1 mol of all of the naturally occurring isotopes of the element

- listed for each element on the periodic table- example 1 mol of iron = 55.85 g/mol

1 mol of zinc = 65.39 g/mol

- some elements exist as molecules such a nitrogen gas

1mol N2 = 2 x 14.01g/mol = 28.02g/mol

Molar Mass of a Compound (M)

- refers to the mass of 1 mol of any pure substance.

- to find the molar mass of a compound use the chemical formula

e.g. CO2 contains 1 carbon and 2 oxygen

1C = 1 x 12.01g/mol = 12.01 2 O = 2 x 16.00g/mol = 32.00

M = 44.01 g/mol

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H2O 2 x 1.01 = 2.021 x 16.00 = 16.00

M = 18.02 g/mol

Ca(OH)2 1 x 40.08 = 40.082 x 16.00 = 32.002 x 1.01 = 2.02

M = 74.10 g/mol

p. 120 #9, 10p. 123 #15, 16, 20p. 125 #3

Calculate the molar mass of the following compounds:

1. PbI2 2. NH4OH 3. CuSO4

4. CaPO4 5. Mn(NO3)5 6. Fe(OH)3

7. NH3 8. S2N4 9. BaSO4

10. C6H12O6 11. NH4HS 12. GaI3

13. CoCl2 14. Cobalt(III) silicate15. Potassium phosphate16. Polonium (II) oxide17. Mercury (II) sulfide18. Fe2(OOCCOO)3

19. Zn(OH)2

20.Cu(NO2)2

21.Co2(Cr2O7)3 22. MgHPO4

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Calculating mass of a sample (m)

Molar mass (M) is equal to the mass of one mole of a compound.

For example the molar mass of water is 18.02 g/mol.

What if we have 2 moles of water?

Then the mass of the water would be 2 x 18.02 = 36.04 g. We use the following formula:

m = nMn = # of moles (mol)m = mass (g)M = molar mass (g/mol)

How many grams are there in 3.5 moles of francium nitride?

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Step 1 – Write the formula and find the molar mass.

Fr3N

3 Fr – 3 x 223.00 = 669.001 N – 1 x 14.01 = 14.01

M = 683.01 g/mol

Step 2 – List what’s given and apply the formula.

n = 3.5 moles m = nMM = 683.01 g/mol = (3.5)(683.01)m = ? m = 2390.54 g

Mass of a substance to moles

If the mass of the sample is given rearrange the formula for “n”

n = m M

e.g.

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How many moles are there in a 16 g sample of carbon dioxide.

CO2 1 x 12.01 = 12.012 x 16.00 = 32.00

M = 44.01 g/mol

m = 16 gM = 44.01 g/moln = ? n = m/M

= 16/44.01

n = 0.36 moles

p. 122 #11-14p. 123 #17,18p. 125 #5,6

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Moles summary

1. Molar mass (M) – must be calculated using the table.

2. Mass (m) – use the formula m = nM

3. Number of moles (n) – use n = m M

p. 135 #29 - 32p. 136 #41 - 43

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Practice Problems Pg 122

Practice Problems Pg123

Check Your Understanding Pg 125

Read Pg 127

Chapter 3 Review Pg 128

Unit 1 Review Pg 134

TEST

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