Download Gr. 11 Review

Gr. 11 Review
Structure of an Atom
Contributions:
The earliest ideas about matter and atoms were developed by Greek
philosophers between 450 and 380 B.C. At that time, the question under
discussion was whether matter had the property of being continuous or
discontinuous.
These concepts can be visualized if you took the "lead" or graphite of
a broken pencil point and divided it in half, then divided that piece in half and
again half of that piece. This process could be continued as long as possible. If
matter was continuous, the process could be continued indefinitely without ever
"running out" of graphite. If matter was discontinuous, then at some point, the
dividing process would end when the last particle which could still be called
graphite was all that remained. A further division would result in destroying the
matter called graphite.


Aristotle: matter is CONTINUOUS
Democritus: matter is DISCONTINUOUS
 Dalton:



Billiard ball model
atom was a solid, indivisible sphere.
Atoms of each element were identical in mass and
their properties.
Atoms of one element differed from that of
another atom
 Thompson:



raisin bun / grapes ‘n jello model
discovered the electron
Positively charged substance filled
the atom. The electrons were arranged
within this substance.
Atom is mostly empty space
 Goldstein:
discovery of proton
 Rutherford:



nuclear atom
Shot alpha particles (+ve) at gold foil, expected
them to go through based on Thompson’s model
Small % were reflected
Hypothesized atom had a postive core – the
nucleus, surrounded by mostly empty space
containing negative charges
 Bohr:
Solar System Model
 Found inconsistencies in
previous models
{Rutherford’s model did not account for the lack of
emission of radiation as electrons move about the
nucleus (at the rate an electron would lose energy,
it would spiral into the nucleus) and the emission of
light at only certain wavelengths (an accelerated electron should radiate
energy at all wavelengths)

Proposed electrons contained specific
energies - referred to as energy levels
The Atom
Particle
Charge
Changing the number of this
particle in an atom means….
Determining #p, #n and #e
#p = atomic number
#n = atomic mass - # p
#e = atomic number
** in a ground state**
Bohr Model Diagram
 Electrons
fill specific energy levels around
the nucleus
 Filling order -
 Determine
the number of protons,
neutrons and electrons of
Neon
p
n
e
 Now
draw the bohr diagram
 Determine
the number of protons,
neutrons and electrons
 Now
draw the bohr diagram
 Determine
the number of protons,
neutrons and electrons for

Na +1

Cl-1

He +2
And draw a Bohr diagram for each.
Schrodinger’s Quantum
mechanical Model

Based on mathematically determining the MOST PROBABLE
LOCATION OF AN ELECTRON



Consists of 4 numbers each giving more information about the
probable location of an electron.
Electrons act like particles and waves
Instead of electrons following distinct paths (Bohr) electron
waves are believed to occupy an orbital space (electron
cloud)
Based on:
Heisenberg’s uncertainty principle:
it is impossible to know both the location and
the momentum of an electron

Shrodinger: Quantum Mechanical Model

Based on determining the MOST PROBABLE LOCATION of an electron

Principle quantum number: (n) specifies the energy level of the electron
 N = 1,2,3,4 to infinity

Azimuthal (Shape) Quantum Number (l): shape of orbital
 l = 0 ‘s’ orbital, spherical
 l=1 ‘p’ orbital – dumbbell
 l = 2 ‘d’ orbital – complex
 l = 3 ‘f’ orbital – complex
 l = 4 ‘g’ and continues through the alphabet

Magnetic Quantum Number (m): spatial orientation - # of orbitals of a
shape
 m = -1 … 0 … +1
 * the NUMBER of ‘m’ values for a given ‘l’ value tells us the NUMBER of
orbitals of that shape (given by the l-value)

Spin Quantum Number (s): determines the spin of the electron
 + 1/2 clockwise
 - 1/2 counter clockwise
In summary
 n = 1,2,3,4 …. to infinity
 L = 0,1,2,3 …. (n-1) 0,1,2,3…. S,p,d,f

m = -1..0..+1
# of m values gives us # of orbitals
as determined by l

S = +1/2 CW
-1/2 ccw
Pauli’s exclusion Principle: NO TWO electrons can
have the SAME set of quantum numbers
THE QUANTUM MODEL
‘l’ value
Shape
n=1
l=0
‘s’ orbital
n=2
l=0
l=1
‘s’ orbital
‘p’
n=3
l=0
l=1
l=2
‘s’ orbital
‘p’
‘d’
m=0
m = -1,0,+1
m = -2,-1, 0, +1, +2
N=4
l=0
l=1
l=2
l=3
‘s’ orbital
‘p’
‘d’
‘f’
m=0
m = -1,0,+1
m = -2,-1, 0, +1, +2
m = -3,-2,-1,0,+1,+2,+3
Major
Level
n
m value
m=0
m=0
m = -1,0,+1
# of sublevels
(different shaped)
n
1 value
1 ‘s’ orb
1 value
3 values
1 ‘s’ orb
1 ‘p’ orb
1 value
3 values
5 values
1 ‘s’ orb
1 ‘p’ orb
1 ‘d’ orb
1 value
3 values
5 values
7 values
# of orbitals
(# of ‘seats’)
n2
** where ‘n’ represents any major level
1 ‘s’ orb
1 ‘p’ orb
n value
1 ‘d’ orb
1 ‘f’ orb
max. # of electrons
2n2
Describe the most probable location of an
electron having the following set of quantum
numbers:
n = 3, l = 2, m = 1, s = +1/2
Describe the most probably location given the
following sets of quantum numbers AND identify the
‘illegal’ set … JUSTIFY your choice(s).
n
l
m
s
a)
2
0
0
+1/2
b)
3
2
-4
+1/2
c)
4
5
+2
-1/2
d)
6
3
-3
-1/2
e)
1
1
0
+1/2
f)
7
5
+3
+1/2
g)
5
3
+4
-1/2
Justification
AUF-BAU PRINCIPLE
 Filing order for electron configuration

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6

Aka spectroscopic notation

Orbital board/box notation: write spec. notation first then
transfer into boxes … watch out for Hunds’ Rule (half –fill
with same spin before doubling up)

** NOTE: a POSTIVE ION is formed by REMOVING electrons
from the OUTERMOST (VALENCE)LEVEL of a NEUTRAL ATOM
… therefore when writing configuration for POSITIVE IONS ..
First write the configuration for the NEUTRAL atom, then
remove electrons from the OUTERMOST LEVEL… the final
configuration may look like it has not follwed proper filling
order .. That is okay (read the first statement in this
paragraph)
Give spectroscopic notation for:
a) P
b)
Fe
c)
Cu+2
d)
Cr+1
 Give
a)
N
b)
Ca
c)
S
orbital board notation for
Give the set of quantum numbers for the
LAST 2 electrons to enter a neutral sulphur
atom (hint: first give spec. notation)
Sketch an outline of a PT
Label/identify:
groups
periods
Metals
Non-metals
Alkali metals
Alkaline earth metals
halogens
noble/inert gases
transition/heavy metals
s, p, and d block elements
-
On your OWN PT transfer the information in the chart on page 7 of your review,
and the reactivity, ionization energy and atomic radius trends
Atomic Radius (AR)
Group
As you go down a group
Atomic Radius increases
(because # of levels increases
∴ valence level is further
away)
Period
As you go across a period
Atomic Radius decreases
(because nucleur charge,
Protons, atomic numberis
increasing creating a
stronger force of attraction
thereby `shrinking`the
atom.. .note: as you go
across a period valence level
is the same)
Ionization Energy(IE): the energy required to
remove the valence (otermost) electron
For a given species .. 1st IE<2nd IE<3rd IE
1st IE – energy required to remove 1st electron
from outermost level
2nd IE – energy required to remove 2nd electron
from outermost level

a)
b)
c)
Given the following ionization energies for
element ‘X’:
What group does ‘x’ belong to?
Give 2 properties of ‘X’
Give the valence and lewis dot diagram for
‘X’
1st IE = 10
2nd IE = 25
3rd IE = 300
4th IE = 380
5th IE = 780
Electronegativity: a measure of electron affinity
(how much a species LOVES electrons)
*does not include noble/inert gases
Man-made scale:0-4,
F is the most EN element
IE/EN trends * identical in the PT
IE/EN are large for small radii
AR IE EN the larger the radius the smaller the IE/EN
AR IE EN vice versa
Group:as you go down a group IE/EN decreases because AR increase
Period: as you across a period IE/EN increases because AR decreases
Comparing anything:
K, Ca, Mg, Na
different groups and periods
1.
Identify valence level for each species
2.
Set up.
small Ar
large IE/EN
large AR
small IE/EN
** second Ionization trends
How to:
1. Rip off one electron (may need to write out
spec. notation)
2. Identify valence level for this set.
3. Do as before.

Ex. Place from low to high 2nd IE
Na, Mg, Li, Cl
Ex. Place from low to high 2nd IE
Na, Mg, Li, Cl
Do questions in review package.
Happiness is … 8 electrons
Octet Rule: a species will react so as to
achieve the electron configuration of the
nearest inert gas
ns2np6
8e- (exception He 1s2)



Metal & non-metal bond
Due to electron transfer (from the metal to the nonmetal)
Attraction of oppositely charged ions
Lewis Dot Diagrams:
Write the lewis dot for each element (# of valence electrons)
b)
Transfer electrons ONE AT A TIME from the metal to the nonmetal, until metal has none and the non-metal has 8 (except for
H…just 2) … bring in atoms of either species if needed
c)
Show the ions (and the number of each) formed
d)
Write the formula
e)
Name the compound .. Metal has the same name, non=metal ion
takes on the ending ‘IDE’
Do EG. 18 on page 9 of review
a)
 Eg.
Na & O


Non-metal/non-metal bond
Due to sharing of PAIRS of electrons so each element will
have 8 electrons on its’ valence level (except H – it will
have 2)
Lewis Dot – indicate the pairs shared
Structural Formula: each pair shared is represented by a dash
It will be easier to draw structural formula first, then replace
each dash with a pair of electrons and then complete the
octet around each element ….
Remember the following rules:
Group
Group
# of bonds
(dashes)
4
5
6
7
4
3
2
1
 Eg.
HCN
 Isomers:
same chemical formula but
different structural formula
 Do
questions 19 and 20, on page 10 of
review
Mole
Avagrados’ #: 6.02 x 1023
Molar Mass (MM): mass of 1 mole of a given substance - g/mole
Molar mass of elements : off PT – eg. Na 23 g/mole
Molar mass of compound: add up all the elements
MM of Mg3(PO4)2 =
% composition
% element A =
MM of A
x 100%
MM of compound
Eg. Mg3(PO4)2
% of Mg.
Formulas:
Mass = # mol x MM
# moles = m / MM
Eg. Det. Mass of 1.6 moles of NaOH
Empirical/Molecular Formula and
Hydrates
 Emprical formula (EF): simplest whole number ratio
 Ie. Sugar is C6H12O6 – EF is CH2O
 Molecular formula (MF): actual formula
 Sugar – C6H12O6
 MF = (EF)n
 n = MM of MF/MM of EF
 Hydrates: compound with a definite # of water
molecules LOOSELY attached to it … salt • X H2O, ratio is
always
 one (1) salt: x (#) water
Lets just get into it!
GAS LAWS
 Boyles Law
 P1V1 = P2V2
 Charles’ Law
 Daltons’ Law: partial pressures, pressure varies
DIRECTLY with # of MOLES of gas particles
 PTOT = Pa + Pb +
Pc
 NTOT = na + nb + nc
 Combined Gas Law:
 ** STP conditions refer ONLY to PRESSURE (101.3KPa) and TEMP (273K)
 Ideal gas law
 PV = nRT
 Lets get right into it – REVIEW QUESTIONS
eg. 20 g of H2 and 30 g of N2 are placed in a 2L container
at 400K.
 A) Determine PTOT
 B) Determine PH2 and PN2
Eg.2Na + 2H20
2NaOH + H2
 What volume of H2 @ STP (T=273K, P= 101.3 Kpa) is
formed when 10 g of Na is reacted?
Do review questions
STOICHIOMETRY – 13 LETTER WORD FOR FUN
REMINDERS
1. BALANCE THE EQUATION!!!
2. Drop arrows
3. Determine your variables given
4. Find what is asked ?
5. FIND MOLES and use mole ratios.
Eg.
2Na + 2H20
2NaOH + H2
What volume of H2 @ STP (T=273K, P= 101.3 Kpa) is formed when 10 g of Na is
reacted?
Review questions
Solution Chemistry
• Arrhenius Theory on ‘ions in solution’
• When a salt (any ionic substance) dissolves in water it breaks up
into its’ ions
• Concentration:
• known as MOLARITY (M)
• moles/litre … moles/L
• moles/L = Molar
• Molarity questions follow stoich
• Formulae:
• M = #moles/V
• Dilution formula: MiVi = MfVf
• Watch for – diluted TO vs diluted WITH
• QUESTIONS
Acid / Base Chemistry
• Properties of
• Acids
Bases
• Arrhenious Definition of:
• Acid: a compound that ionizes (dissociates) IN SOLUTION to
form (give) ____________________
• Base: an ionic _____________________ that dissociates IN
SOLUTION to produce (give) ______________________ ions
• ** presence of H+1/OH-1 ions IN SOLUTION give an Acid/Base
its’ properties**
•
• Man made scale 0 – 14
• 0 -6.9 =
7=
7.1-14 =