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 lives
- WHMIS informs workers about the chemicals in three ways
- 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.
- MSDS (Material Safety Data Sheet) was required to accompany every chemical bought and sold (p.7).
Classifying Matter
Matter is defined as anything that has mass and volume
Matter may be solid, liquid or gas.
Matter
MixturePure Substances
Combinations of matter Matter that has a
that can be separateddefinite composition.
by physical means
Do not have a
definite composition
Gravel Milk GoldWater (H2O)
Assignment:
- State whether the following is a pure substance or a mixture:
a) sea waterb) ironc) bronze
d) 14k golde) table saltf) oxygen
- State whether the following mixtures are homogeneousorheterogeneous:
a)Oil-and-vinegar salad dressing
b)Steel
c)Cranberry juice
d)Sugar dissolve in water
e)Milk
f)Antifreeze
- State whether the following pure substances is an element or a compound:
a) Copperb) Water, H2O
c) Methane, CH4d) Silver
- Classify each of the following substances:
a) Graphite, Cb) Shampoo
c) Coffeed) Motor oil
- Name the following elements:
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, C
Hf, 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)
Developing (History of) Atomic Theories
- 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
- J.J. Thomson (1856-1940)
-English physicist
-Atoms contain negatively charged electrons
-Electrons are like raisins in a plum pudding or “raisin bun” model.
- Ernest Rutherford (1871-1937)
-Atom contains electrons and positively-charged particles
-Atom composed of
- A nucleus – a central region that is positively charged and contains most of the mass
- protons are heavy positive particles within the nucleus
- Electrons – particles with a negative charge and are very light (compared to protons).
- Electrons circle around the nucleus
- Empty space surrounding the nucleus is very large within which electrons move (planetary model).
- Rutherford also proposed existence of the neutron to account for the mass difference between hydrogen and helium
- Neutrons are heavy particles like protons but have no charge
- Isotopes are atoms of the same element that differ in mass (but are chemically alike). (element with different number of neutrons)
- Niels Bohr (1885-1962)
-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).
- 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
-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
for iron and sodium.
Fe Atomic # = 26
Atomic mass = 55.85
p = 26
e = 26
n = 56 – 26 = 30
Note: when finding the number of neutrons we round the atomic mass to the nearest whole number.
Na Atomic # = 11
Mass # = 22.99
p = 11
e = 11
n = 23 – 11 = 12
Assignment:
Find the number of protons, electrons, and neutrons of the elements with atomic numbers 1 to 30 and 40-70.
p. 11 #1-3 (copy questions or complete sentences)
Nuclear Notation Continued
- 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 = 6
Atomic number#n = 12 – 6 = 6
13 6C#n = 13 – 6 = 7
- Another notation used is: e.g. Lithium–7 or Li - 7
where 7 is equal to the atomic mass.
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 8
p.24 # 1 to 4
p. 37 (b, c) (copy question first or complete sentences)
p. 38 - define the key terms (first column)
Bohr’s Model of the Atom
According to Bohr’s modelelectrons 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 1storbit can hold a max. of 2 electrons
The 2nd a max of 8
3rd 8
4th 18
5th18
6th32
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 must
have 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 number
It’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
We can draw the orbits using a simplified version.
e.g. Cobalt
___ 9
___ 8
___ 8
___ 2e
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)
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 Metals
Group 2 – Alkaline Earths
Group 17 – Halogens
Group 18 – Noble gases or Inert gases
-‘Staircase’ line separates metals from non-metals
-Metalloids border this line
Francium is the most reactive metal.
Fluorineis themost reactive non-metal.
Valence Electrons
-The outermost occupied energy level (orbit) of an atom is called its valence energy level.
-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 beLi
e.g. Oxygen
#p = 8
#e = 8
#n = 8 ____ 6
____ 2e
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 Friday
Assignment:
Draw electron dot diagrams for:
- Scandium
- Fluorine
- Beryllium
- Vanadium
- Gallium
p. 38 – Define the key terms and
p. 38 # 7, 8, 13
Quiz
Draw electron dot diagrams for:
- Sc
- Na
- Chromium
- Ar
Ions
-Ions are atoms, which have gained or lost electrons, in order to become more stable – it happens during chemicalreactions.
-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-)
- 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-metalnon-metal
Assignment
p.36 #1-3, 5
p.38 #6, 9, 15,16 (copy question or complete sentences)
Compounds
There are two basic types:
- Ionic
- 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.
In ionic compounds electrons are traded.
e.g. NaCl (p.30 – sketch Fig. 1.27 here)
Molecular Compounds
-Atoms which shareelectrons 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.
Molecularcompounds – 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)
Investigation 1-A Pg 33
Check Your Understanding Pg 36
Read Pg 37
Review Pg 38
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
- The first element in the compound is the one most left on the periodic table.
- The suffix ‘ide’ is attached to the name of the second element.
- Prefixes are used to indicate how many atoms of each type are present in one molecule of the compound.
Prefixes:
1 = mono6 = hexa
2 = di7 = hepta
3 = tri8 = octaMEMORIZE
4 = tetra9 = nona
5 = penta10 = 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:
- Silicon dioxide
- Sulphur monoxide
- OF2
- SiBr4
- PH3
- N2O
- CO
- NBr3
- P2I3
- SO3
- N2O4
- Tetraphosphorous hexaoxide
- Dinitrogen tetraoxide
- Heptasilicon monobromide
- Octaboron decaiodide
- B2O3
- BrF7
- N3O6
- H2Cl5
- Triselenium diastatide
- Diarsenic pentaoxide
- Sulphur trioxide
- C3O2
- C2H6
- As3Br7
- SO2
- Selenium monoxide
- Diboron trioxide
- PF3
- P2O5
- P4O10
- Arsenic trifluoride
- BrF7
- Hydrogen chloride
- 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
- The first element in the name of the formula is the metal
- The second element, the non-metal, is named as an ion. The suffix ‘ide’must be present.
- No prefixes are used.
e.g.
Fe2O3 – Iron oxide
CuS – Copper sulfide
KCl – Potassium chloride
p. 45 #5 - 7
p. 46 #9, 10
p. 47 #12
p. 48 #3, 5 (copy question)
Writing Formulas for Binary Ionic Compounds
In an ionic compound the total number ofpositive chargesmustequalthe total negativecharges – the compound must be electrically neutral.
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+
Cl1-
Therefore the formula is MgCl2
Chromium oxideCr3+O2-
Cr3+O2-to balance the charges we use a shortcut method – charges are “traded” across.
Cr2O3
Calcium oxide Ca2+O2- Ca2O2CaO
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
Copper (II) oxideCu2+O2- CuO
Tin (IV) fluorideSn4+F1- SnF4
PbI2 Lead (II) iodide
Pb2+ I1-
Cr2S3Chromium (III) sulfide
Cr3+S2-
Is this formula correct LiOLi1+O2-
No – correct formula is Li2O
p. 47 #11
p. 48 #4, 5
p. 49 #7, 9
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 ammoniumion, 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, bracketsmust 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
NH4NO3 Ammonium nitrate
Al(NO3)3 Aluminum nitrate
Sodium sulfateNa1+(SO4)2-Na2SO4
Na1+
SO42-
Na1+
Ammonium phosphate(NH4)1+(PO4)3-
(NH4)3PO4
Gallium hydrogen carbonate
Ga3+(HCO3)1-Ga(HCO3)3
Assignment:
Practice Problems p. 52 #13-16
Practice 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
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, 5
p. 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.
WaterH2OorHOH
OzoneO3
AmmoniaNH3
Hydrogen PeroxideH2O2
MethanolCH3OH
EthanolC2H5OH
SucroseC12H22O11
GlucoseC6H12O6
MethaneCH4
Diatomic and Polyatomic Elements
If these elements are FREE, that is ALL ALONE, they are written as:
H2P4
N2S8
O2
F2MEMORIZE
Cl2
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.
Assignment:
p. 79 (a, b, c, e,)
p. 80 – key terms
p. 81 #25
Test tomorrow.
Properties of Acids and Bases
- 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 acid
H2SO4-Sulphuric acid
- 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
0acid 7base 14
p.80 #12,13
p.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
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 Sulfurousacid
HClO2 Hydrogen chlorite Chlorous acid
p. 70 #20 (a, b, c) #21 (a, b, c)
#22,23
p. 71 #6
p. 79 (e, k) and p. 135 #26-28
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
-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
-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 following
- energy change – heat and/or light
- exothermic – release energy
- endothermic – absorb energy
- odour change
- colour change
- formation of a gas – bubbling
- formation of a solid – precipitate
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.
PbI2(s)
NH4OH(aq)CuSO4(aq)
p.90 #1, 2
p. 93 #1- 4
p.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