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UNIT 8Atomic Theory
KEY IDEASVocabulary What does it mean?
bright line spectrum
ground state
excited state
photon
quantum
orbital
electron configuration
shell
subshell
core notation
valence electron
atomic radius
ionization energy
electronegativity
covalent bond
polar covalent bond
dipole
ionic bond
electrostatic attraction
Lewis structure
octet rule
resonance
Vocabulary What does it mean?
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bond energy
bond length
proton
electron
neutron
ion
anion
cation
atomic number
mass number
isotope
Atoms of an element with a different number of _____________.
atomic mass
The average mass of a particular element (includes all isotopes) can be calculated by
atomic mass = (%A)(atomic massA)+ (%B)(atomic massB)+ (%C)(atomic massC)+ ……
8.0 – Science 10 Review (P. 146 #13-19, P.150 #23)
Review: Remember the subatomic particles in an atom.
Subatomic Particle Symbol Location Mass Chargeproton
neutronelectron
Atomic number: equals to ____________________Mass number: ___________________Atomic mass: ___________________________
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Example: Calculate the atomic weight of carbon, given the information below.
Practice: Calculate the atomic weight of Zn, given the information below.
Nuclear Notation (P. 149 #22): There are two ways to describe atoms:
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X – mass number Xmass number
atomic number
Nuclear Symbol
charge
Isotope Abundance in Nature (%) Mass (amu)
carbon-11 0 11.011carbon-12 98.9 12.000carbon-13 1.1 13.003carbon-14 0 14.003
Isotope Abundance in Nature (%) Mass (amu)
Zn-64 48.6 63.93Zn-66 27.9 65.93Zn-67 4.1 66.93Zn-68 18.8 67.92Zn-70 0.6 69.93
Example:
Write the symbol for a neutral Fluorine-19 atom __________
Write the symbol for a Barium-137 atom that has had 2 electrons removed __________
Write the symbol for neutral Carbon-13 __________
Example: Circle the isotopes of C613 from the list below.
C614 C7
13 C612 C8
11 C611
8.1 – Bohr’s Model
Recall Bohr’s model of the atom:
Bohr’s model of the atom was famous because it could explain the _____________________________________. Recall that a prism splits light into its individual components.
The spectrum of white light is _________________________, it shows all the colours of the rainbow
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Energy level (shell) max. # electrons
1st
2nd
3rd
4th
The spectrum of energized gas is _______________________, it shows ________________________.When particles are exposed to energy
________________ (lowest energy state) atoms absorb specific energy and become “___________”o an electron ___________________________________
excited atoms then release the energy at specific frequencies to get back to ground stateo the electron moves back to its original shell, emitting a _________________________________
the difference between two particular energy levels is called a _________________ the lines in the spectrum are produced when electrons de-excite
Bohr’s model was later proven significantly wrong in 2 ways:
_____________________________________________________________________________________
_____________________________________________________________________________________
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8.2 – Quantum Mechanics: Schrödinger’s Atomic Model
Schrödinger’s model describes the _______________________ of where to find an electron in an atom
___________: the ____________________ around a nucleus where an electron can be found
orbitals are described by ________________________
Quantum Number Symbols What does it Mean?
1st or principal quantum number (n)
2nd quantum number (l):
3rd quantum number (ml):
4th quantum number (ms):
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Each energy level has a specific set of orbitals and each one represents where a maximum of ______ electrons can be found.
Orbital Type Begins at n= # of Orbitals in a Subshell
Maximum # of Electrons in Subshell
8.3 – The Energy Level Diagram
Instead of representing atoms with the Bohr diagram, we can represent them with the more accurate energy level diagram.
Rules to follow when filling orbitals:1. Fill orbitals from lowest to highest energy (Aufbau Principle)2. Place one electron in each orbital of a sub-shell3. When each orbital of a sub-shell has one electron, go back and pair the electrons (Hund’s Rule)4. If two electrons are in a orbital, they must have opposite spin (Pauli Exclusion Principle)
These rules ensure that the electron configuration gives the ___________ energy, most __________ atom by reducing _______________________________________.
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Example: Fill the orbitals with He electrons
How many electrons are in He? ____
How many shells have electrons? ____
How many sub-shells have electrons? ____
How many orbitals have a single electron? ____
How many orbitals have paired electrons? ____
What is the electron configuration?
Example: Fill the orbitals with Na electrons
How many electrons are in Na? ____
How many shells have electrons? ____
How many sub-shells have electrons? ____
How many orbitals have a single electron? ____
How many orbitals have paired electrons? ____
What is the electron configuration?
Example: Fill the orbitals with Al electrons
How many electrons are in Al? ____
How many shells have electrons? ____
How many sub-shells have electrons? ____
How many orbitals have a single electron? ____
How many orbitals have paired electrons? ____
What is the electron configuration?
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8.4 – Electron Configuration
The periodic table is a tool to obtain the electron configuration of elements quickly.
Simply write the orbitals as they appear on the period table above, going left to right, row by row. Each element space counts as one electron.
Example: Write the electron configuration for the following atoms.
a) Ar
b) Ga
c) Ag
d) Rn
Core notation: It sure is annoying to write super long electron configurations! the shortcut: look for the closest previous noble gas element to the element you are writing the
configuration for and start there useful because we aren’t interested in the core electrons anyways (they don’t participate in chemical
reactions) to write core notation for a noble gas, use the previous noble gas
Example: Write the electron configuration for Ga using core notation.
Closest previous noble gas element: ______ Core notation: _______________________
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Practice: Write the electron configuration in core notation for the following elements.
a) Zn
b) K
c) Kr
There are 2 exceptions to the electron configuration of elements up to Kr:
What you would expect In Reality
Cr
Cu
Reason: ______ filled and ______________________ filled subshells are more stable than partially filled subshells, and in Cr and Cu, it only takes one electron to achieve this stability in the 3d subshell.
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8.5 – Valence Electrons (P.29 #29)
When atoms undergo a chemical reaction, only the outermost electrons are involved (furthest from nucleus and highest in energy) _________________
valence electrons are the ones in the _______ and _______ subshells beyond the noble gas core, and any partially filled d and f subshells
if all subshells are completely full in a atom (i.e. noble gas) then there are no valence electrons
Example: How many valence electrons are in the following atoms?
Atom Electron Configuration # Valence e-
AlGaPbXe
Can you see a pattern for counting valence electrons on the periodic table?
8.6 – Electron Configuration of Ions
For negative ions, add the extra electron(s) where you left off in the neutral atom
For positive ions, electrons in subshells with the ___________________ value and highest energy are removed first electrons are removed from the _____-subshell first, then ______, and then ______-subshells
Example: Write the electron configuration for the following ions:
Neutral Atom Ion
a) O2- →
b) S2- →
c) Br- →
d) Sn2+ →
e) Sn4+ →
f) V2+ →
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When an atom becomes an ion, its electron configuration sometimes becomes the same as the configuration for the nearest _______________________. When two atoms have the same electron configuration, they are said to be __________________________.
Practice: Which of the following is isoelectronic with krpton when ionized?
Cl Zn Br Ca Sr Se I
Consider this: If two atoms have the same electron configuration, are they the same element?
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UNIT 8 PART 2The Periodic Table & Trends
Review (P. 164 #35-39)
8.7 – Atomic Radii
Atomic radius: the size of an atom.
Across a period, the atomic radius ____________________ because the _____________________________o
o
Down a group, the atomic radius ____________________ because the _____________________________o
o the number of electrons also increases, ___________________________________________
Example: Consider the following pairs of atoms. Which atom has the larger atomic radius?
a) O and O2- ______
b) Ca and Ca2+? ______
Example: Arrange the following species in order of increasing size
Rb+ Y3+ Br- Kr Sr2+ Se2-
Species # p # e- Order of Increasing SizeRb+
Y3+
Br-
KrSr2+
Se2-
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Consider This: Which has a larger effect on atomic radii, a change in the number of protons, or a change in the number of electrons?
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8.8 – Ionization Energy (IE) (P. 168 #48-51, P. 169 #52)
Ionization energy: the energy needed to __________________________ from a neutral atom the harder it is to remove the e-, the __________________ the IE
Consider this: Explain the trend seen in the first ionization energy for the first 20 elements on the periodic table.
1. Gen era lly
speaking, what happens to IE1 as you go across a period and why?
2. What explains the high IE1 for noble gases?
3. How does IE2 compare with IE1?
4. What happens to IE as you go down a group and why?
5. What explains the decrease in IE1 from Be to B and Mg to Al?
8.9 – Electronegativity
Electronegativity: an atom’s ability to ____________________________
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Atomic # Element IE1 (kJ/mol) IE2 (kJ/mol)1 H 1312 -2 He 22372 55103 Li 519 72854 Be 900 17565 B 799 24276 C 1088 23537 N 1406 28558 O 1314 33849 F 1682 3372
10 Ne 2080 396211 Na 498 456112 Mg 736 145013 Al 577 181714 Si 787 157715 P 1063 190316 S 1000 225117 Cl 1255 229718 Ar 1519 268919 K 418 306720 Ca 590 1145
moving across a period: increaseso
moving down a group: decreaseso
Summary of Trends
Review P. 170 #53-55
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UNIT 8 PART 3Bonding & Structures
8.10 – Types of Bonds
Chemical bond: force that holds atoms together to form molecules, the attraction between electrons of one atom to the nucleus of another atom
The nature of the bond (degree of electron sharing) can be predicted using _____________________
Covalent bonds occur between 2 non-metals that have an electronegativity difference of ________________ equal sharing of electrons a pair of electrons is simultaneously attracted to two nuclei two atoms share the bonding electrons atoms will try to get full outer shells
Example: hydrogen gas H2
How many bonds are there? ______ ( __________________________)
Difference in electronegativity between the two atoms: _____________________
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Example: oxygen gas O2
How many bonds are there? _______ ( __________________________)
Difference in electronegativity between the two atoms: _____________________
Polar covalent bonds occur between 2 non-metals that have an electronegativity difference of ______________ unequal sharing of electrons a pair of electrons is shared, but they are more attracted to one nuclei than the other, resulting in a
________________
Example: HCl
Let’s just look at the pair of electrons that are shared between H and Cl
Electrons are more attracted to Cl because Cl has a higher electronegativity
H becomes __________________________________ and Cl becomes __________________________________
Example: H2O
Again, let’s just look at the pair of electrons that are shared between H and O
O has higher electronegativity than H
H becomes __________________________________ and O becomes __________________________________
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Ionic bonds occur between a metal and non-metal that have an electronegativity difference of ______________ no electron sharing, involves electron transfer
o metal atom (or the atom with the lower electronegativity) loses electrons and becomes positively charged ___________
o non-metal atom (or the atom with the higher electronegativity) gains electrons and becomes negatively charged ____________
atoms are held together by ___________________________________ (attraction of opposite charges)
Example: NaF
Difference in electronegativity between the two atoms: __________________
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8.11 – Lewis Structures
Lewis structures (electron dot structures) describe valence electron arrangement around atoms of a molecule
Drawing atoms element symbol represents nucleus + core electrons dots represent valence electrons dots are drawn in pairs as a reminder that electrons are paired in orbitals
Example: Draw the Lewis structure for the following atoms
Li Be B C N O F Ne
Drawing Covalent Molecules
octet rule predicts bonding arrangement: bonded non-metallic atoms have 8 electrons in their outermost energy levels (exception ______ can only have 2)
molecules tend to be symmetrical covalent and polar covalent bonds are represented by pairs of dots between two atoms number of dots you draw must equal the sum of the valence electrons of all atoms in the molecule pairs of electrons forming covalent bonds can be represented by a line
Example: Draw the Lewis structure for the following molecules
a) CCl4 b) NH3
c) C2H6 d) CO2
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Example: Draw the Lewis structure of HCN
There are some exceptions to the octet rule besides H there are elements that don’t have enough valence electrons to achieve a full octet
(________________) there are elements that can accommodate more than an octet (_____________________________)
a) BeI2 b) BF3
c) PCl5 d) SF6
Drawing Ionic Compounds use square brackets to indicate that the atoms are held together by electrostatic charge
Example: Draw the Lewis structure for NaCl
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8.12 – Formal Charges
Sometimes molecules can be drawn in more than one way. The question is which one is the real one?
Example: Nitrous acid HNO2
To determine which structure is the best and most stable, we use formal charges as a check. formal charges indicate how many electrons are “owned” by an atom
the sum of all formal charges in a molecule must equal to the molecule’s overall charge the most stable molecule is one where the formal charges are the ________________
Example: Draw the Lewis structure for sulfuric acid.
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8.13 – Charged Molecules and Resonance Structures
To draw the Lewis structure for charged molecules the number of electrons drawn needs to take the charge into account enclose the molecule in square brackets and put the charge in superscript
Example: Draw the Lewis structure of NH4+
Sometimes molecules have fractional bonds as a result of electron movement among bonds to accurately represent the molecule, show all possible structures each of these structures is called a _______________________________
Example: Draw the Lewis structure for SO32-
The real molecule is all of the above structures, with each of the O atoms having 113 bonds, therefore when you
try to represent SO32-, you must draw all three structures.
Example: Draw the Lewis structure for HPO32-
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8.14 – Bond Energy & Bond Length
Bond energy: the measure of __________________________________ the stronger the bond, the ________________ the bond energy expressed as heat required to break 1 mole of molecules
It takes energy to _________ bonds, and energy is released when bonds are _____________
Example:
a) The reaction turning 1 molecule of H2(g) into two H(g) atoms involves bond _________________. making / breaking
Energy is ___________________. released / required
b) The reaction turning two oxygen atoms into a molecule of O2 involves bond _________________. making / breaking
Energy is ____________________. released / required
c) Compare the bond energies of single, double, and triple bonds. Explain the data.
d) Compare the bond energies of C-N, C-O, C-F. Explain the data.
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Bond length: the distance between 2 nuclei of a molecule the shorter the bond length, the ______________ the bond, the ______________ the bond energy
Example: How do the bond lengths of single, double, and triple bonds compare with each other?
For ionic compounds, two things must be considered in order to determine bond strength. They are (in order or importance):
1. The charge: the _____________ the charge on each ion in the ionic compound, the ________________ the ionic bond, the ______________ the bond energy
2. Electronegativity: the higher the electronegativity of the non-metal, the _______________ the attraction, the ______________ the bond energy.
Example: Rank the following compounds in order of increasing melting point (increasing bond strength).
LiF LiCl LiBr
Example: Which of the following ionic compounds has the higher bond energy?
NaCl MgO
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