Download Review sheet Atomic Structure, Electron Configuration, and Periodic

Unit 4 Packet - Page 1 of 9
Honors Chemistry – Unit 4
Chapters 5 & 6: The Periodic Table and Electron Configuration
Vocabulary: Monday, Feb 14
Element Trading Cards: Thurs. Feb. 10
Vocabulary:
Element
Noble gas
Alkali metal
Periodic Law
OBJECTIVES:
Chapters 6
Periods
Nonmetal
Actinide
De Broglie
Problem Set Due: Monday, Feb 14
Test Date: Wednesday, Feb. 16
Families
Metal
Metalloid
Group
Lanthanide
Valence
Schrodinger
Cannizzaro Mendeleev
Hund
Halogen
Transition metal
Moseley
Aufbau
Alkali Earth Metals
Heisenberg
Identify groups as vertical columns on the periodic table.
Know that main group elements in the same group have similar properties, the same number of
valence electrons and the same oxidation number.
Understand that reactivity increases as you go down within a group for metals and decreases for
nonmetals.
Identify periods as horizontal rows on the periodic table.
Identify representative (main group) elements as groups 1, 2, 13-18.
Identify alkali metals, alkaline earth metals, halogens, and noble gases based on location on periodic
table.
Identify transition elements as groups 3-12.
Use a periodic table to write the symbols of elements, given their names.
List the characteristics that distinguish metals, nonmetals, and metalloids.
Assign the oxidation number for elements and each element in the formula of a chemical compound.
Define valence electrons and state how many are present in atoms of each main-group element.
Locate and name the four blocks of the periodic table and explain the reasons for these names.
Discuss the relationship between group configurations for the atoms of any element using orbital
notation, electron-configuration notation, and when appropriate, noble-gas notation.
Understand how the periodic table is set up.
Be able to list and describe the four quantum numbers.
Be able to write electron configurations using the periodic table.
Be able to explain the uncertainty principle, Hund’s rule and the exclusion principle.
Be able to write Noble gas configurations, orbital diagrams and dot diagrams.
Understand the wave-particle duality of nature.
Be able to define above vocabulary
Given a set of at least 4 elements, be able to place them in order of increasing/decreasing periodic
properties.
Be able to order ions and atoms by size.
Be able to identify the families on the periodic table.
Unit 4 Packet - Page 2 of 9
Unit 1 PROBLEM SET:
Review Problems:
1. Calculate the mass of a sample of copper that occupies 4.2 x 10-3 cm3 if the density of gold is 8.94
g/cm3.
2. How much aluminum phosphate can be produced from the reaction of 2.8 g of aluminum nitrate with
3.3 g of lithium phosphate?
3. If 4.04 g of N combine with 11.46 g O to produce a compound with a molecular mass of 108.0 what is
the molecular formula of this compound?
Current Unit Problems: (Electron Config)
4. Identify the group of elements that corresponds to each of the following generalized electron
configurations:
A. [noble gas] ns2np5
B. [noble gas] ns2(n-1)d10np1
C. [noble gas] ns2(n-1)d2
5. Why does the element molybdenum have an electron configuration ending in 5s14d5 rather than
5s24d4? Explain your reasoning.
6. Explain why the He atom has a smaller radius than the H atom.
7. A. How many protons, electrons and neutrons in Ba2+ and Cl-1
B. Write the electron configurations for Ba2+ and Cl-1
8. Write the complete electron configuration for Uranium.
Current Unit Problems: (Periodic Table)
1. In an experiment each student needs to weigh out 1.84 g of Copper (Cu) wire. If we only have a spool
of wire that weighs 50.0 g, how many students can do the experiment?
2. What property did Mendeleev use to organize his periodic table?
3. Which of these sets of elements have similar physical and chemical properties?
a. oxygen, nitrogen, carbon, boron
b. strontium, magnesium, calcium, beryllium
c. nitrogen, neon, nickel, niobium
4. Give 3 properties of metals.
5. Give 3 properties of nonmetals.
6. Give 3 properties of metalloids.
7. What is so special about the noble gases?
8. Complete Element Symbol worksheet.
9. Complete Periodic Table Puzzle.
10. Complete Periodic Table worksheet.
11. Complete Atomic Structure worksheet.
Unit 4 Packet - Page 3 of 9
Chapter 4 - The Arrangement of Electrons in Atoms
Reread section 4-1 (we covered it in unit two). Answer the questions below as you read section 4-2 pages
98-104
1.
2.
3.
4.
5.
6.
What did DeBroglie hypothesize about the behavior of electrons?
What is the Heisenberg uncertainty principle?
How did Schrodinger contribute to the development of the idea of atomic orbitals?
What is the quantum theory?
What is an orbital?
What are quantum numbers?
Chapter 5 - The Periodic Table
Answer the following questions as you read pgs 123-127
1. How did the following chemists contribute to the development of the periodic table?
A. Cannizzaro
B. Mendeleev
C. Moseley
D. Ramsay
2. What does the term periodic mean?
3. What is the periodic law?
**Be able to locate the following on the periodic table**:
alkali metals, alkaline earth metals, transition metals, halogens, chalcogens, noble gases, lanthanides
and actinides
QUANTUM NUMBERS AND ATOMIC ORBITALS
FOUR QUANTUM NUMBERS:
Quantum #(Q#) = #s that specify the properties of atomic orbitals and their e-s.
There are 4 Quantum #s
1. Principal Quantum Number: (n)
 Indicates energy levels (shells), n = 1, 2, 3 etc. Diagram =
 As n increase so does energy
 # e-s on a level = 2n2 Example: e-s on energy level 2?
2. Orbital Quantum Number: ( )
 indicates the shape of an orbital and the subshells or sublevels

subshells: s, p, d, f

n = 1 , energy level one has 1 subshell named 1 s

n = 2, has 2 subshells named 2s, 2p

n= 3 has

n= 4 has
subshells named
each subshells contains orbitals (think of them as the electron’s home)
s has 1 orbital, p has 3 orbitals, d has 5 orbitals, f has 7 orbitals.
subshells named
Each orbital can hold a maximum of 2 e-s.

Therefore:
s = 1 orb,
2
e-
p=
orbs,
e-
d=
orbs,
e-
f=
orbs,
e-
Unit 4 Packet - Page 4 of 9
Shapes of orbitals:
3. Magnetic Quantum Number ( ): indicates orbital orientation about the nucleus.
See example on board:
4. Spin Quantum number: ( ):
indicates spin of e2 values = +1/2 and – ½ (clockwise or counterclockwise)
Pauli’s exclusion principle:
Homework/Classwork: Fill in following chart!
Fill in the following Tables:
Quantum Number
Name & Symbol
Significance
1
2
3
4
Energy Level
Subshells
# of Orbitals
Maximum
number of e-s
Total number
of e-s
n=1
n=2
n=3
n=4
***Electron Configuration Notes (you will to use your own paper to take notes on this
section)***
Unit 4 Packet - Page 5 of 9
Valence Electrons
Unit 4 Packet - Page 6 of 9
Lewis Dot Diagrams
Unit 4 Packet - Page 7 of 9
CHAPTER 5 - PERIODIC PROPERTIES
Define words with star (*) using pgs140-154. As a class we will fill in the other blanks.
I.
Atomic Radius * Most important Group Trend: size __________________as you go down a group; b/c of the increase in
energy levels
Period Trend: size _______________ as you go across the period from left to right; Due to the fact that
as you go across the pull of the nucleus increases.
Example on board Li and F.
Together: Put in order by increasing size:
Li
O
Si
Ar
and N
With partner: Place in increasing size:
K
II.
P
As
Na
F
Ionization energy(first) * 2nd ionization energy*
Period Trend (most important): I E _________________________ across the periods from left to right.
Metals tend to lose e- ‘s = low IE
Nonmetals tend to gain e- ‘s = high IE
Group Trend: IE increases _____________ the group
Atoms with less e-‘s hold more tightly onto the e-‘s they have.
Together: Put the following in decreasing order of ionization energy
F
Cl
Mg
Sr
P
With partner: Put the following in order of increasing IE
Na
III.
Li
Te
Se
Electron Affinity * Period Trends (most imp): e- affinity increases across the table from _____ to __
Group Trends: electron affinity increases _______ the group.
(Trend is the same as trend for IE)
Practice: Place the following in order of decreasing e- affinity
Mn
IV.
Br
S
Rb
Fe
Electronegativity *:
Developed by Linus Pauling
(Do not include the Noble Gases – Why?)
Period Trends (most imp)- increases across the period
______________ least electronegative ________________ most electronegative
Group Trend electonegativity increases _______ a group.
Most electronegative element:__________
Least Electronegative element:_____
Practice: Place in order of increasing electronegativity:
W
Bi
I
Rb
As
Unit 4 Packet - Page 8 of 9
V.
Ionic Radius – tells us the size of the ions
Cations = ______ ions
Anions= _____ ions
Cations: ___ e-s (therefore have + charge) they are smaller than their parent atom
Ex: Na+ is smaller than Na
Anions: _______e-s (have ________ charge) they are larger than their parent atom. Ex: Cl-1 is larger
than Cl
Practice: Which is larger?
Ca or Ca2+
Al3+ or Al
N or N3VI
C or C4+
Reactivity:
Metals: most reactive ___________ and _______________
Nonmetals more reactive ____________ and _______________
(Except Noble gases- Why?)
Practice:
Circle the most reactive metal:
Au
Cs
Sr
Ti
Circle the most reactive nonmetal:
Se
Cl
O
Ne
Decreases across a period
Increases down a group
Increases across a period
Decreases down a group
Increases across a period
Decreases down a group
Review sheet Atomic Structure, Electron Configuration, and Periodic Properties
1. What is the total number of electrons needed to fill the fourth energy level?
2. Identify the following rules/scientists:
A. an electron occupies the lowest energy orbital that will receive it
B. electrons fill empty orbitals before pairing up
C. no 2 electrons in the same atom can have the same 4 quantum numbers
P
Unit 4 Packet - Page 9 of 9
D. It impossible to know both the exact speed and position of an electron
3. Write the long form and noble gas notation of As
4. Draw the dot diagram for As
5. Given: C Mg S Ba place these elements in order of:
A. increasing atomic radius
C. increasing electron affinity
B. decreasing electronegativity
D. decreasing ionization energy
6. Write the noble gas configuration for Bi
7. Explain the charge on a sulfide ion using the concept of electron configurations.
Answers:
1. 2n2 = 32
2. A. Aufbau principle B. Hund’s rule C. Paul’s exclusion principle D. Heisenberg’s uncertainty principle
3. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3 ;
[Ar]4s2 3d10 4p3
4. As with 1 pair of dots and 3 simgle dots
5. A. C S Mg Ba
B. S C Mg Ba
C. Ba Mg C S
D. S C Mg Ba
6. [Xe] 6s2 5d10 4f14 6p3
7. Sulfide’s charge is 2-, this is because sulfur’s ending configuration is 3s2 3p4;to become stable suflur needs to gain
2 more electrons (to make an octet = 8 e-s which is very stable); if sulfur gains 2 e-s its ending configuration becomes
3s2 3p6 matching the very stable noble gas Ar’s configuration