Honors Chemistry - Topic 3 Test (Sections II and III) Outline

Honors Chemistry - Topic 3 Test (Sections II and III) Outline

Atoms - The Building Blocks of Matter

II. Bohr Model of The Atom

pp. 303-312

A. The Nature of Light [5.2a, 4.1a end “Bohr”; 5.2b]

1. Types of spectra

·  Continuous spectrum – entire or most of electromagnetic spectrum

·  Line Spectrum – pieces/little bits of color seen for element

2. Concepts of light

a. light as a wave

·  See SG from short quiz

b. light as a particle - photons

·  Atoms that have gained extra energy release that energy in the form of light.

o  Line spectrum: very specific wavelengths of light that atoms give off or gain

o  Each element has its own line spectrum, which can be used to identify that element.

o  The line spectrum must be related to energy transitions in the atom

•The atom is quantized, i.e. only certain energies are allowed.

c. electromagnetic spectrum

3. Quantizing line spectra: Rydberg Equation

o  used in atomic physics to describe the wavelengths of spectral lines of many chemical elements.

· 

B. The Bohr Atomic Model

1. Explaining the existence of line spectra

·  The Bohr model was based on a simple postulate, Bohr applied to the hydrogen atom the concept that the electron can exist only in certain energy levels without an energy change but that, when the electron changes its state, it must absorb or emit the exact amount of energy that will bring it from the initial state to the final state.

·  (The ground state is the lowest energy state available to the electron. The excuted state is any level higher than the ground state. )

·  The formula for a change in energy (∆E) is:

o  ∆Eelectron = Efinal – Einitial

2. Bohr’s Electron configurations

·  Energy of atom is related to distance of electron from nucleus

3. Bohr’s Energy-level diagram and quantum numbers

o  Energy of the atom is quantized

§  Atom can only have certain specific energy states called quantum levels or energy levels.

§  When atom gains energy, electron “moves” to a higher quantum level

§  When atom loses energy, electron “moves” to a lower energy level

§  Lines in spectrum correspond to the difference in energy between levels

o 

§  Ground state: minimum energy of an atom

·  Therefore electrons do not crash into the nucleus

§  The ground state of hydrogen corresponds to having its one electron in the n=1 level

§  Excited states: energy levels higher than the ground state

o  Distances between energy levels decrease as the energy increases

§  1st energy level can hold 2e-1, the 2nd 8e-1, the 3rd 18e-1, etc.

§  Further from nucleus = more space = less repulsion

o  Valence shell: the highest-energy occupied ground state orbit

Problems with Bohr Model

o  Only explains hydrogen atom spectrum (and other 1-electron systems)

o  Neglects interactions between electrons

o  Assumes circular or elliptical orbits for electrons (which is not true)

III. Modern Atomic Structure – Wave Mechanical Model

Wave mechanical - a new model of hydrogen atom that seemed to apply equally well to all other atoms, which Bohr’s model failed to do.

A. Changing Bohr’s Model [4.1b] p.312-318; 319-327; 352

·  Orbital Theory – Electrons are found in orbitals around the nucleus. The orbitals describe regions of space where an electron is most likely to be found at any given moment.

o  s-sublevel – shape of sphere around nucleus. Electron is located somewhere within this space.

·  Quantum Mechanical Model (aka Charge Cloud Model) - treats electrons as waves and uses wave mathematics to calculate probability densities of finding the electron in a particular region in the atom

o  –Schrödinger Wave Equation

o  –Can only be solved for simple systems, but approximated for others

· 

1. Duality of matter: DeBroglie (Wave Mechanical Model of the Atom)

o  Suggested matter consist of waves

o  Experiments later showed that electrons could be treated as waves

o  Just as light energy could be treated as particles

2. Heisenberg’s Uncertainty Principle

·  Heisenberg- Says you can’t know where electron is and its is going at the same time (b/c any effort to determine one is going to change it).

o  Heisenberg Uncertainty Principle – can’t actually describe where an electron is in an atom. All you can do is define an area of space around the nucleus where you are most likely to find an electron (Led to orbital theory)

3. Matter-Waves: Schrödinger’s wave concept

·  Schrödinger – Came up with equation that describes the motion of a single electron around a single proton of hydrogen

B. Describing The Modern Atom

As we mentioned earlier, each principal energy level, n, has n sublevels. This means the first has one sublevel, the second has two, the third has three, etc. The sublevels are named s, p, d, and f.

Energy level principal quantum number, n / Number of sublevels / Names of sublevels
1 / 1 / s
2 / 2 / s, p
3 / 3 / s, p, d
4 / 4 / s, p, d, f

At each additional sublevel, the number of available orbitals is increased by two: s = 1, p = 3, d = 5, f = 7, and as we stated above, each orbital can hold only two electrons, which must be of opposite spin. So s holds 2, p holds 6 (2 electrons times the number of orbitals, which for the p sublevel is equal to 3), d holds 10, and f holds 14.

GOOD TABLE BELOW

Sublevel / s / p / d / f
Number of orbitals / 1 / 3 / 5 / 7
Maximum number of electrons / 2 / 6 / 10 / 14
Quantum number, l / 0 / 1 / 2 / 3

1. Electron configurations [5.3a] – pg.75

It is important to remember that, when there is more than one orbital at a particular energy level, such as three p orbitals or five d orbitals, only one electron will fill each orbital until each had one electron.

·  (Each electron has a charge of -1)

·  Each period represents the principle energy level

·  Principle Energy levels are made of Sublevels

o  There are 4 basic sublevel types

§  s-sublevel – can hold up to 2 electrons

§  p-sublevel - can hold up to 6 electrons

§  d-sublevel – can hold up to 10 electrons

§  f-sublevel - can hold up to 14 electrons

o  (# sublevels = principle energy level #)

a. Aufbau Principle – the principle that an electron occupies the lowest energy orbital that can receive it.

b. Hund’s Rule(of Maxium Muliplicity) – principle that, after this(above/Augbau), pairing will occur with the addition of one more electron to each orbital.

c. electron blocks and the periodic table

·  s, p, d, and f blocks

·  Box Diagrams – show from most general (energy levels) to most specific (orbitals)

o  each sublevel is made of orbitals

o  1 box=1 orbital

o  Orbital – region of space where you are likely to find 2 electrons, which have opposite spins

§  s-sublevel – 1 orbital (box)

§  p-sublevel - 3 orbital (boxes)

§  d-sublevel – 5 orbitals (boxes)

§  f-sublevel – 7 orbitals (boxes)

o  Spins are represented by arrows (up and down) – up first, followed by down

This matches the size of the s-block.
The p subshell can hold up to six electrons. This matches the size of the p-block.
The d subshell can hold up to 10 electrons. This matches the size of the d-block with 10 columns.
The f subshell can hold up to 14 electrons. The 14 columns of the f-block match the filling of the f subshell.

· 

d. Lewis dot structures [5.3b] –Page 78

·  In 1916, G.N. Lewis devised the electron dot notation , which may be used in place of the electron configuration notation. The electron dot noation shows only the chemical symbol surrounded by dots to represent the electrons in the incomplete outer level (valence electrons).

·  Lewis Dot Structures – 2 dots on each side of letter (maximum amount is 8)

o  Valence electrons – electrons in outermost energy level (most possible is 8 = stable octet)

§  Family 1A 1 valence electron

§  Family 2A - 2 valence electrons

§  Family 3A - 3 valence electrons

§  Family 4A - 4 valence electrons

§  Family 5A- 5 valence electrons

§  Family 6A- 6 valence electrons

§  Family 7A - 7 valence electrons

§  Family 8A – 8 valence electrons

o  Examples – # dots=# of valence electrons

§  Sodium (3s) has 1 valence electron and dot diagram would have a single dot

§  Magnesium (3s2) has 2 valence electrons and has two dots

§  Aluminum has 3 valence electrons has 3 dots

§  Tin (Sn) has 4 valence electrons and 4 dots

§  Nitrogen has 5 valence electrons and 5 dots

§  Oxygen has 6 valence electrons and 6 dots

§  Bromine has 7 valence electrons and 7 dots

§  Argon has 8 valence electrons and 8 dots (stable octet of valence electrons)

2. Quantum numbers [5.4a] (Electron Zipcode/General Location is space)

·  There are 4 quantum numbers that recognize a specific electron in orbital notation (See more info below)

· 

1.  n -

§  Ex. 1st principle energy level has a quantum number of 1, 2nd PEL has quantum number of 2, 3, 4, etc.

2.  ℓ - Sublevel (ℓ) – different sublevels are given dif. quantum #’s

ℓ / Sublevel
0 / s
1 / p
2 / d
3 / f

3.  m ℓ - Orbital (m ℓ) – which orbital the electron is located in

§  0 is central orbital

4.  ms - Spin (ms) – Box diagram: Up arrow (+ ½), down arrow (- ½)

·  You can name an electron with 4 quantum numbers, and find a specific electron if given the 4 numbers.

n / The principal quantum number – only know this name
l / The subsidiary or azimuthal or angular momentum or orbital shape quantum number
ml / The magnetic quantum number
ms / The electron spin quantum number

3. Atomic Orbitals

a. shapes of orbitals = s, p, d and f

·  Another method is commonly used to designate sublevels. Instead of giving the quantum number ℓ, we use a letter (s,p, d, or f) to indicate the sublevel. A sublevel for which ℓ=0 is refered to as an s sublevel. If ℓ=1, we are dealing with a p sublevel. A d sublevel is one for which ℓ=2, in an f sublevel, ℓ=3

o  A sublevel ______

b. Pauli Exclusion Principle

·  States that in a given atom no two electrons can have the same set of four quantum numbers (n, l, m1 and ms)

o  No orbital may have more than 2 electrons.

o  Electrons in the same orbital must have opposite spins.

o  s sublevel holds 2 electrons (1 orbital)

o  p sublevel holds 6 electrons (3 orbitals)

o  d sublevel holds 10 electrons (5 orbitals)

o  f sublevel holds 14 electrons (7 orbitals)

4. Excited and ground state [5.4b; 5.4c]

·  The ground state is the lowest energy state available to the electron.

·  The excited state is any level higher than the ground state.

o  Ex. 1s22s233s1 is excited state of Boron b/c skipped over 2p

o  If matches periodic table, it is the ground state

o  If electrons are in a level they are occupied, if they reach maximum amount, then they are full

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