chem - ch. 1 Notes Teacher

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 WHMIS helps you tell if a material is hazardous, type of risk, how to handle a
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1
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FAMILY/GROUP: elements in vertical columns – have similar properties.
Reactivity INCREASES from
top to bottom within groups
1 and 2
Reactivity DECREASES from top
to bottom in groups 17 & 18.
(Fluorine is the most reactive
non-metal)
 PERIOD: elements in
horizontal rows
– these elements change
from metals to nonmetals.
 Alkali Metals:
 Group 1 metals
 Soft, silver coloured metals
that react violently with H O
2
to form basic solutions
 Most reactive: cesium &
francium
2
 Alkaline Earths:
 Group 2A metals
 Light, reactive metals that form oxide coatings when exposed to air
 Halogens:
 Group 17, non-metals, highly reactive.
 Fluorine is the most reactive
 Noble Gases:

Group 18

Generally unreactive (inert)
Reading the periodic table:
3
Significant digits:
 1. All digits 1-9 are significant
 Ex. 123 ( 3 sig digs)
With zeros it gets more complicated…
 3. leading zero’s are not significant
 Ex. 0.12 (2 sig digs)
 4. Zero’s to the right are only significant if:
 A. units are attached to the number
 Ex. 100 g (3 sig digs)
4
B. In scientific notation 1.0 x 10 (2 sig digs)
c. placed after other digits but behind a decimal point are significant; 7.90 (3
sig digs)
Complete: WHMIS Worsksheet
1.1
Structure of Matter
The Atom; smallest part of an element
Is made up with 3 major particles:
 Protons (positive charge particle)
 Electrons (negative charge particle)
 Neutrons (neutral particle)
4
Atomic number – the number of protons in an element
Calculating electrons:
 Atomic number in element
Calculating neutrons:
Atomic mass = n + p
Neutrons = atomic mass – protons
Always remember to subtract!
Example:
# of protons
Element?
# electrons
# of neutrons
Mg
Fe
Br
Pb
More examples:
Element
# of protons
(Atomic #)
# of electrons
# of neutrons
Lithium (Li)
Titanium (Ti)
Carbon (C)
Chlorine (Cl)
5
Bohr Model:
 Niels Bohr developed a diagram to represent electrons and
the energy levels of an atom.
 Each energy level can hold a certain maximum number of
electrons equal to the number of elements in the period
Bohr Diagram Steps:
Step 1: see what row the element is in,
and count how many elements in
the row it is. (include element)
All rows before the last one, are
fully occupied with the max
amount of electrons
Include protons and neutrons at the bottom
Examples:
Oxygen Atom
Sodium Atom
6
Lewis Dot Diagram – deals only with valence electrons
Valence electrons – electrons in the last energy level
Step 1: write symbol
Step 2: determine valence electrons
Step 3: place electrons in North, East, South and West positions. Place one
at a time, with two maximum at each position.
Carbon Example:
Examples:
Lithium
Calcium
Neon
helium
fluorine
** Atoms have one mission: to gain, loose, or share electrons, in order to become
fully satisfied.
7
 Diatomic: H , O Share pair of electrons
2
2
 Polyatomic: share many pairs of electrons
ex. Ammonium
Atoms:
 Most stable when its outer energy level is full of electrons.
 A maximum of 8 electrons may occupy a valence energy level.
This is known as the octet rule.
 If an atom gains electrons, it become negative charged
 If it loses electrons, it become positively charged
 Anion: element ion that gained electrons (negatively charged)
 Cation: electron ion that lost electrons (positively charged)
Properties of:
 Metals: conduct, “magnetic”, malleable, ductile, loose e-, cations, silver,
solid (except Hg)
 Non-metals: brittle, dull, does not conduct, gains e-, anions
Can a metal and non metal bond?
 Yes, metal donates e- to the non metal (Ionic Bond)
What about calcium oxide?
All compounds have a net charge of 0, thus neutral (stable)
8
Complete: Page 8: 1, 2, 3; Page 9: 5; Page 13: 7; Page 14: 1, 2, 3, 6; Page 21: 11;
Page 23: 1, 2, 4
1.3 Breaking Bonds
Chemical change – occurs when a new substance is formed, accompanied by a
change in color, odor, state, or energy
 Exothermic: chemical change when energy, usually in the
form of heat, is released into the surroundings.
 Endothermic: a chemical change in which energy is absorbed
from the surroundings. (substance usually gets colder)
Solutions
 Solute: a substance that dissolves (salt)
 Solvent: a substance that does the dissolving and is in greater
proportion in the mixture. (H 0)
2
Solutions are a good medium for chemical change, water is a good solvent
9
Why Water is a Good Solvent?
 All 3 physical states occur on earth under “normal” conditions
 Boiling point and melting point is higher than similar substances
 Polar molecule – molecule with a partial positive and
negative charge
 Electrostatics – attracted to objects that are positive and/or
negative (water will bend toward a charged object)
 Universal solvent
o dissolve ionic compounds
o + and – charges are produced with in sol’n
 Dissociation- separation of an ionic compound into individual ions in a
solution
o smaller parts (compound  ions)
 Ionic compound: has a cation (metal) and an anion (non metal)
 Molecular compound: 2 non metals
Electrolytes:
 electrolytes (conducts electricity when dissolved water)
 non-electrolytes (does not conduct electricity when dissolved in
water)
10
Non- electrolytes:
 “Like dissolves like”
◦ Water can not break down covalent bonds in molecular compounds
Concentration:
 amount of solute to solution (ratio)
Concentration Vs Dilution
qualitative analysis (observable property)
◦ color
◦ taste
◦ scent
◦ test the solution
 Conductivity
 pH
(A greater concentration results in a higher/more intense property)
11
quantitative analysis (see pg. 10 in dada book)
 C V =C V
i
i
f
f
C= molar concentration
V= volume (L)
 C= n/V
i= initial solution
F=final solution
n= # of moles (mol)
Diluting Solutions:
Example:
CV = C V
i i
f f
You have 65.0mL of a 0.759 mol/L solution of sodium chloride, NaCl
(aq).
Calculate the final concentration of the solution if it is diluted to a final volume of
100.0 mL.
12
Example:
CV = C V
i i
f f
You have 100.0mL of a 1.490mol/L solution of sodium chloride, NaCl
(aq).
Calculate the final volume of the solution if it is diluted to a final concentration of
0.500 mol/L
 Assignment: Page 27: 16, 17, 19; Page 30: 21, 23; Page 31: 24a-c; Page 33:
1-4; Page 39: 1, 2, 4, 5
13
1.5 CALCULATING CONCENTRATION / DILUTIONS
 % V/V = percent by volume concentration
 (V solute / V solution) x 100
 Most commonly used for liquids dissolved in liquids
 PPM = parts per million
6
 (m solute/ m solution) x 10 ppm
 Used for very dilute solutions
 ALWAYS … SOLUTE / SOLVENT!
Helpful hints:
 Read the question carefully
 Is the solute included in the solvent?
 Remember the significant digits rule
 Remember the conversion factor for mL to L
 1L = 1000mL
 Know how to read powers of ten in your calculator
 See page 43
Steps for completing problems:
 Step 1: write down what you know and don’t know
 Step 2: write down and rearrange your formula
 Step 3: plug in your numbers and include units
14
Percent by volume example:
 A hair product requires you to combine 20.0 mL of hydrogen peroxide with
enough water to produce a solution with a total volume of 120.0 mL.
Determine the percent by volume concentration of the solution.
PPM example:
-3
A 200 g sample from a bottle of water contains 5.4x10 g of mercury.
Calculate the concentration of mercury in the sample in parts per million.
15
Concentration example:
A sample of water taken from a nearby lake is found to have 0.0035 mol of salt in
a 100 mL solution.
Determine the concentration of the salt in the lake.
ACIDS:
 a substance that produces hydrogen ions when dissolved in water to form a
conducting aqueous solution.
 Turns blue litmus paper red.
 Conduct electricity (they are electrolytes)
 Taste sour
 Their pH ranges from 0 to 7
*Always add an acid to water, never add water to the acid.
Bases:






Bases are empirically defined as a substance that:
Taste bitter
Conduct electricity (they are electrolytes)
Feel soapy or slippery
Turns red litmus blue
Their pH ranges from 7 to 14
Assignment: Page 42: 28-30; Page 44: 31-33; Page 46: 34-36; Page 50: 3738; Page 52: 1, 3, 4, 6a,b, 7a
16
Dilutions:
Standard solutions Are:
 Are precise
 For comparisons
 Volume
 Color
Preparing a solution:
 Example prepare 500 ml of a 1.2 mol/L copper II sulfate solution. And then
dilute it to prepare 100 ml of a 0.12 mol/L
 calculate the required mass necessary to prepare the solution using n =cv
and m= nM
Solution 1: v1= 500mL  0.500L C1= 1.2 mol/L
Solution 2: v2= 100mL  0.100L C2= 0.12mol/L
n=cxv  0.500L x 1.2mol/L = 0.60mol
M = nxM  0.6 mol x 159.62 g/mol = 95.8g
(CuSO4) = 63.55+ 32.07+ (4x16) = 159.62g/mol
17
 weigh the mass out on a electric scale
 pour the mass in to a 500 ml beaker
 dissolve the mass with approximately ¼ the required volume of solution
 transfer the solution to a 500 ml volumetric flask using a funnel
 add enough water to fill the volumetric flask
 I then need to calculate what volume of copper II sulfate solution is
required to make my diluted solution
 n=cv of new solution
 v=n/c of old solution
 New diluted solution:
 n= c xv  0.12mol/L x 0.100l = 0.012mol
2
2
 V= n/c  0.012mol/1.2mol/L = 0.01L = 10mL
1
 pipet out that volume from my original solution and transfer it to a second
100 ml volumetric flask
 raise volume to 100 ml
18
Diluting solutions:
CiVi = CfVf
You have 85.0mL of a 0.655 mol/L solution of sodium chloride, NaCl(aq).
Calculate the final concentration of the solution if it is diluted to a final volume of
150.0 mL.
19
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