4.1

Electron Configurations and Orbitals • In reality, electrons travel in more complicated paths than those shown by Bohr models. In order to show this, we write "addresses" for electrons known as electron configurations. configurations Example:

2

2

1s 2s 2p

Energy Level (Shell)

6

# of e -

Sub-shells (shape) The diagram at left shows the paths traveled by the 10 electrons whose electron config. is shown above. 1s (1st shell/sphere) sphere) 2s (2nd shell/sphere) sphere)

2p (2nd shell, dumb-bells) dumb ~1~

• s - Orbitals: available to the atom in all shells. -

spherical shape 1 pair e- / 1 orbital path max capacity: 2 e- / s2 1s, 2s, 3s, 4s, 5s, 6s, 7s ....

• p - Orbitals: available starting with the 2nd shell. -

dumb-bell or figure-8 shape 3 pairs e- / 3 orbital paths max capacity: 6 e- / p6 2p, 3p, 4p, 5p, 6p, 7p

• d - Orbitals: available starting in the 3rd shell. - clover-leaf shape - 5 pairs e- / 5 orbital paths - max capacity: 10 e- / d10 - 3d, 4d, 5d, 6d ... (the picture above shows ONE d-orbital, not all 5) ~2~

• f - Orbitals: available starting in the 4th shell. - complex, multi-dimensional dimensional - 7 pairs e- / 7 orbital paths - max capacity: 14 e- / f14 - 4f, 5f...

(the picture above shows ONE of 7 ff-orbitals) •

So, how does this relate to the Bohr Model? 1st Shell: 2 e-

2nd Shell: 8 e-

3rd Shell: 18 e-

4th Shell: 32 e~3~

• Larger atoms build more complex orbital shapes to accomodate more electrons. New orbitals build through, around, or on top of prior ones. Example:

Nucleus

1s2 2s2 2p6 3s2 3p6

1s2

1s22s2

1s22s22p63s2

1s22s22p6

1s22s22p63s23p6

• The periodic table is built to show this order. ~4~

• The 4 blocks of the periodic table:

• Period #s are equal to the shell #s with 2 exceptions: d-block block is always 1 behind (period 4 = 3d) f-block block is always 2 behind (period 6 = 4f)

~5~

• To write an electron configuration, you "read" the periodic table from left to right and write completed addresses until you reach the element you seek. If you "stop" part way through a block, you only give that part of the address. Examples:

Oxygen = 1s22s22p4 Cobalt = 1s22s22p63s23p64s23d7

~6~

Practice: give the "final" address in each of the following element's electron configurations. Ex.1) Carbon, C

2p2

Ex.6) Scandium, Sc

Ex.2) Phosphorus, P

Ex.7) Argon, Ar

Ex.3) Iridium, Ir

Ex.8) Cerium, Ce

Ex.4) Lithium, Li

Ex.9) Aluminum, Al

Ex.5) Plutonium, Pu

Ex.10) Platinum, Pt

Give the full electron configuration for each: Example: Nitrogen = 1s22s22p3 Ex.11) Be

Ex.14) F

Ex.12) P

Ex.15) B

Ex.13) Cr

Ex.16) He ~7~

Remember the octet rule? Non-metals complete a NEW octet by filling s and p-orbitals in the same energy level (shell) Ex. 2s22p6 or 3s23p6 Ex.17) Oxygen, O

1s22s22p4 → 1s22s22p6

=

Ex.18) Nitrogen, N

1s22s22p3 →

=

Ex.19) Fluorine, F

→

=

Ex.20) Carbon, C

→

=

O2-

Metals "shed" their excess electrons rather than attempting to construct new octets. Ex.21) Aluminum, Al 1s22s22p63s23p1 → 1s22s22p6 = Al3+ Ex.22) Magnesium, Mg 1s22s22p63s2 →

=

Ex.23) Lithium, Li

→

=

Ex.24) Beryllium, Be

→

=

~8~

For ions, add e- to anions and take e- away from cations. Empty shells are not written. Ex.25) Na+

1s22s22p63s1

→

Ex.26) O2-

1s22s22p4

→

1s22s22p6

Ex.27) Cu2+ Ex.28) HAn easier way to write e. configurations is called a Noble Gas Configuration and only writes the configuration past the last noble gas. Ex.29) Iodine, I

[Kr] 5s24d105p5

Ex.30) Praseodymium, Pr Ex.31) Uranium (II) cation, U2+ Ex.32) Chloride anion, Cl~9~

Using subshells and orbitals, draw quantum mechanical models of the following atoms. Ex.33) Neon, Ne: 1s2 2s2 2p6

Ex.34) Iron, Fe: 1s22s22p63s23p64s23d6

~ 10 ~

4.2

Orbital Diagrams Orbital diagrams are a visual way of drawing electron configurations and follow these rules: Aufbau Principle: "An electron must occupy the lowest available energy orbital" - configurations are built from the lowest energy orbital to the highest. Ex. 1s → 2s → 2p → 3s ....... Hund's Rule: Fill orbitals at the same energy level with one electron before pairing them. (Remember, electrons repel / don't "like" each other.)

Ex.

2p

NOT

2p

Pauli Exclusion Principle: Two electrons cannot occupy the same "address" in a configuration. In an orbital, electrons must "spin" in opposite directions. Possible:

Impossible: ~ 11 ~

Ground State vs. Excited Electrons: normally, electrons will occupy the lowest energy level possible within an orbital, the "ground state". Heat, electricity, and light can give energy to an electron causing it to become "excited" and jump into a higher orbital or shell. + Energy

2p3 Ground State

2p2 3s1 "Excited" or Promoted

Normally, we assume atoms are in the ground state. When excited electrons return to the ground state, they emit photons of light. Practice: Give the full electron configuration for these atoms and draw their orbital diagrams: Ex.1) Helium, He

Ex.2) Beryllium, Be

~ 12 ~

Larger atoms may contain p, d, and even f-type subshells. According to Hund's Rule, we fill the orbitals within the subshells with SINGLE electrons before pairing them. Use noble gas configurations to draw each of the following: Ex.3) Sulfur, S

Ex.4) Neon, Ne

Ex.5) Fluorine, F

Ex.6) Fluoride ion, F-

~ 13 ~

Much larger atoms will require careful use of Aufbau (5s before 4d, for example), Hund, AND Pauli. Write noble gas configs and orbital diagrams for each: Ex.7) As

Ex.8) Cf3+

Ex.9) Te-

Ex.10) K+

~ 14 ~

Evaluate each of the following orbital diagrams. If incorrect, which principle(s) or rule(s) is violated? Ex.11) 3p

4s

Ex.12) 4s

3d

Ex.13) 2s

2p

1s

2p

Ex.14) 3s

Ex.15) 6s

4f

~ 15 ~

4.3

Types and Properties of Chemical Bonds A chemical bond is formed when two or more atoms give, take, or share valence electrons. The difference in electronegativity (∆χ or "delta chi" ) between two elements determines the nature of the bond which will be formed. 0 / Small ∆χ

Medium ∆χ

Large ∆χ

e- are shared / e- shared unequally / e- given/taken 2 Non-Metals, Similar χ: Non-Polar Covalent Bonds Electrons are shared equally or 50:50 2 Non-Metals, Differing χ: Polar Covalent Bonds Electrons are shared unequally, not 50:50 Metal(s) and Non-Metal(s): Ionic Bonds Electrons are given and taken. Pure Metals or Alloys: Metallic Bonds Electrons are mobile, able to move freely. ~ 16 ~

Covalent Bonds result from the sharing of electrons between two non-metals in a sort of "tug of war". A single particle of a covalent compound is called a molecule, i.e. a single H2O. When atoms are similar in electronegativity, the sharing is mostly equal and non-polar covalent. Look for non-metal combinations where both are highly electronegative or both are weak. F - F

C - H

C - N

S - Cl

Polar covalent bonds form between highly electronegative non-metals and weaker non-metals. This unequal sharing creates areas of partial positive and negative within the bond. O - H

H - Cl

P - F

Si - Br ~ 17 ~

Ionic Bonds occur when the difference in electronegative strength is so great (metal/non) that one atom steals electrons from another. A single ionic particle, i.e. a single NaCl, would be called a formula unit.

Na

Cl

K

S

K

Na

Cl

K

S

K

Ca

O

Ca

O

Even though no electrons are shared, ionics are bonded by their opposing charges, a concept known as electrostatics. ~ 18 ~

Metallic Bonds have unique properties due to the fact that electrons move freely between atoms. This effect creates a "sea of electrons" surrounding the positive nuclei of the metal atoms. +

+

+

+

+

+

+ +

+ +

+ +

In reality, there is no such thing as a single metallic bond. There is a general attraction in all directions between metal nuclei + and the sea of e- ( ) allowing metals to be strong, yet still flexible: Metallic bonds are.....

+

Conductive (mobile e-) Malleable / Ductile Dense ~ 19 ~

Ex.1) NiBr3

Bond: Ionic

Particle: Electrons? Formula Units Give/Take

Ex.2) Cu

Metallic

Atoms

Ex.3) NO2

N-P Cov.

Molecules

Fair Share

Ex.4) H2CO3

P. Cov.

Molecules

Unfair Share

Ex.5) Au + Cu

Metallic

None, Alloy

Sea of e-

Ex.6) CHCl3 Ex.7) FeCl3 Ex.8) AuCl4 Ex.9) CS2 Ex.10) Fe2O3 Ex.11) HF ~ 20 ~

Sea of e-

Bond:

Particle:

Ex.12) ArF2 Ex.13) Fe + Mn Ex.14) NO3 Ex.15) Cr + Fe Ex.16) RhO4 Ex.17) IrI4 Ex.18) OBr2 Ex.19) NaF Ex.20) Pure Zn Ex.21) Zn + Cu Ex.22) O2 ~ 21 ~

Electrons?

4.4

Lewis Dot Diagrams A Lewis Dot Diagram for a single atom or ion is created by writing the symbol of an element surrounded by "dots" showing valence electrons. Start from the top and work clockwise.

Ex.1)

Oxygen =

O

Lithium =

Li

Aluminum =

Al

Nitrogen = N

Ions will have more or less, depending on charge Ex.2)

O2-

(6 VE → 8 VE)

O

Al+

(3 VE → 2 VE)

Al

Practice: Draw a lewis dot diagram for each ion. Ex.3) Mg+ =

P2-

=

Mg

C3- =

C

P

Na+ =

Na

~ 22 ~

Lewis Dot for compounds, these rules apply: #1 Calculate the total number of VE. #2 Place the least powerful atom in the center, surrounded by stronger (more electronegative) elements. Typically, this is the "single" element in the compound. (Never group 1, always C) Examples:

#3 Covalent: Ionic:

chlorine in HClO4 nitrogen in NH4+

oxygen in K2O carbon in CO2

Bond all atoms to center using single bonds (solid line, each requires 2 e-) Skip to step 4, no bonds (no lines).

#4 "Feed" the remaining electrons to elements in order of electronegative strength to satisfy the octet rule. #5 If COVALENT and OUT of electrons, "borrow" existing pairs of electrons to create double and triple bonds and satisfy the octet rule. ~ 23 ~

Covalent bonds are formed by the sharing of one or more electron pairs. These pairs are represented by a solid line in Lewis Dot diagrams: Ex.4) OF2

VE: Bonds: Electrons: Particle:

Ex.5) CCl2Br2

VE: Bonds: Electrons: Particle:

Remember, hydrogen is stable with 2 e- : Ex.6) SiH4

VE: Bonds: Electrons: Particle:

~ 24 ~

Cations and anions must be shown with brackets [] and charges when drawn as a Lewis Dot Diagram. Ex.7) H3O+

VE: Bonds: Electrons: Particle:

Ex.8) NH2-

VE: Bonds: Electrons: Particle:

Ex.9) ClO4-

VE: Bonds: Electrons: Particle:

Ex.10) OH-

VE: Bonds: Electrons: Particle: ~ 25 ~

Covalent compounds may also contain double or triple bonds. This occurs when there aren't enough electrons to satisfy all individual atoms and additional sharing must occur: Ex.11) CO2

VE: Bonds: Electrons: Particle:

Ex.12) HCN

VE: Bonds: Electrons: Particle:

Some Lewis diagrams with double/triple bonds can be drawn multiple ways. This is called resonance: Ex.13) CO32-

~ 26 ~

Ionic compounds involve a give and take of electrons and are bonded through their opposing charges. This concept is known as electrostatics, and is often simplified to "opposites attract". Survival of the strongest (most electronegative), no sharing here! Ex.14) Li3P

VE: Bonds: Electrons: Particle:

Ex.15) CaO

VE: Bonds: Electrons: Particle:

Ex.16) AlCl3

VE: Bonds: Electrons: Particle:

~ 27 ~

4.5

VSEPR and Molecular Geometry As you know, chemical bonds are made from electrons (given, taken, shared, etc). Consequently, bonding electrons repel (push away) from one another as well as NON-bonding electrons.

H2O Formula

HOH Lewis Diagram

( AX2E2 / Bent)

This repulsion creates the shapes (Molecular Geometry) of molecules. Overall this idea is called VSEPR or Valence Shell Electron Pair Repulsion. We use AXE notation to match molecular geometries to Lewis Dot diagrams: • A = the central atom (Usually only one) • X = the number of bonded atoms (not bonds) • E = the number of unbonded pairs of e- around the center atom. ~ 28 ~

We can match a substance's AXE notation with specific 3D shapes or molecular geometries: :

AX

AX2

AX4

AX3E

AX3

AX5

AX2E & AX2E

AX6

Hint: you will want to memorize the shapes above. Practice by using the molecular modeling kits in lab (ask me), the PHET simulator, and making flashcards. This will be quizzed relentlessly in class and lab! ~ 29 ~

Given each of the following Lewis Dot diagrams, determine the AXE notation for the central atom of each and identify the corresponding molecular geometry:

Ex.1)

Ex.2)

Ex.3)

Ex.4)

~ 30 ~

You can also use your knowledge of VSEPR to identify the molecular geometries of central atoms within larger molecules: C1 = Ex.5) C2 =

C3 =

C1 =

Ex.6)

C2 =

O3 = ~ 31 ~

For each of the following substances: construct a Lewis Dot diagram, determine the AXE, and identify the matching molecular geometry: Ex.7) Methane, CH4

Ex.8) Nitrate ion, NO3-

Ex.9) Boron trichloride, BCl3

Ex.10) Carbon disulfide, CS2

~ 32 ~

(Boron is stable w/ 6 VE)

Ex.11) Sulfur hexafluoride, SF6

Ex.12) Phosphorous pentafluoride, PF5

Ex.13) Phosphate ion, PO43-

Ex.14) Silicon oxyfluoride, SiOF2

~ 33 ~

4.6

Intermolecular Forces (IMFs) IMFs are the forces which exist between molecules, as opposed to bonds which exist within molecules. Water, for example, has both bonds and IMFs:

Bonds

IMF In biology, you learned about a specific IMF known as "hydrogen hydrogen bonding". bonding While not a chemical bond, bond hydrogen bonds are quite powerful IMFs and produce the liquid state and surface tension present in water. Why is this important? IMFs give compounds their unique properties, such as state and solubility. solubility ~ 34 ~

The "give and take" nature of ionic bonds creates fully-positive cations and fully-negative anions. anions The forces of attraction between these ions causes them to form a repeating 3D pattern known as a crystal lattice.. Table salt, NaCl, for example:

Na

Cl

Na

Cl

Cl

Na

Cl

Na

Na

Cl

Na

Cl

For ionics, the bonds ARE the IMFs Opposites attract or "electrostatics".

The strong IMF/bond pairing in ionics causes them to all exist as crystalline solids at room temperature. ~ 35 ~

Metallic bonds are similar to ionics in that the bond and the IMF are the same. "Sea of Electrons"

+ Unlike ionics, metallic bonding is omni-directional. This means that metals atoms are NOT rigidly locked in place. This gives metals strength AND flexibility: malleability and ductility. Voltage Heat

+ +

+

When voltage is applied, the electrons "flow" through the metal, creating electricity. Those same electrons carry the vibrations from heating, explaining thermal conductivity. ~ 36 ~

Non-polar covalent bonds are fairly weak in terms of IMFs. CH4 (methane....farts) for example: H H I I H ─ C ─ H H ─ C ─ H I I H H The bonds between carbon and hydrogen are nonpolar and produce no significant + or - charges. As a result, small non-polar compounds have minimal IMFs and are gases at room temperature. Larger non-polars, such as oils, are massive enough to generate a new IMF: London Dispersion Forces (essentially "flickers" in the electron cloud). This allows them to be liquids or solids at room temperature. Oil, tar, and other massive hydrocarbons are examples. ~ 37 ~

Polar bonds can produce very strong IMFs under the right conditions. Hydrogen bonding in water, for example. Look for molecules with clear asymmetric positive and negative regions. Chloroform, CHCl3, is a liquid due to the attraction between partial positive and negative regions in the molecule: H I Cl ─ C ─ Cl I Cl

Cl I Cl ─ C ─ Cl I H

Radial symmetry, however, cancels out this effect: F I F ─ C ─ F I F

F I F ─ C ─ F I F

Carbon tetrafluoride, CF4, is a gas at room temp. ~ 38 ~

Intermolecular Forces (IMFs)

Ionic: Electrostatics Rigid crystal lattice Brittle solids (All!) Full + and e-

-

e

e-

e-

e-

e-

+ + + + + + -

e-

e

e

e-

-

ee-

Metallic: Strong IMF "Sea of electrons" Malleable, ductile Solids (exc. Hg) Conductors Polar Molecules Medium IMF Partial +/Solid OR liquid Non-Polar Polar Molecule Weak IMF London Disp Forces. Small = (g) / Big = (l,s)

~ 39 ~

Draw Lewis dot diagrams for each of these substances. Predict the strength of IMFs and state (solid, liquid, gas) for each of these compounds. Ex.1) NH3 (ammonia)

(Lewis Dot Diagram)

Ex.2) CO2

(Lewis Dot Diagram)

~ 40 ~

Valence e- : Element types: Bond(s): Electrons: Particle: IMFs: Predicted State(s):

Valence e- : Element types: Bond(s): Electrons: Particle: IMFs: Predicted State(s):

Ex.3) CaF2

(Lewis Dot Diagram) Ex.4) N2

(Lewis Dot Diagram)

Valence e- : Element types: Bond(s): Electrons: Particle: IMFs: Predicted State(s): Valence e- : Element types: Bond(s): Electrons: Particle: IMFs: Predicted State(s): Valence e- : Element types: Bond(s): Electrons: IMFs:

Ex.5) C8H18 (octane) (No Lewis Dot) Particle: Predicted State(s): ~ 41 ~

Ex.6) HF

(Lewis Dot Diagram) Ex.7) MgO

(Lewis Dot Diagram)

Valence e- : Element types: Bond(s): Electrons: Particle: IMFs: Predicted State(s): Valence e- : Element types: Bond(s): Electrons: Particle: IMFs: Predicted State(s): Valence e- : Element types: Bond(s): Electrons: IMFs:

Ex.8) CH3OH (ethanol) (No Lewis Dot) Particle: Predicted State(s): ~ 42 ~

~ 43 ~