Test 3 Flashcards by Benjamin Pratt

4th quantum spin(s)

Electron is a spinning charge, can either spin up or down. Represented by either +1/2 or -1/2

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Auf bau principal

lowest energy occupied first

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Polyexclusion principal

each electron has a unique set of quantum numbers

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Hund’s principal

parallel spins in different boxes before pairing up to minimize C-e repulsion

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Magnetic properties

-related to unpaired electrons

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Diamagnetic

Unaffected by magnetic field, all electrons paired

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Paramagnetic

Attracted by magnetic field, has unpaired electrons

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Electron configuration of Ions

Cations loose e
Highest n always looses first
If more than one highest n highest l lost first

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Electron configuration of Anions

gain e to get to noble gas configuration

add e as usual

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Valence electrons

part of s and p orbitals and non filled f and d orbitals

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non valence electrons

filled d and f orbitals are not valence electrons

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Electron configuration of Cr

[Ar}4s^13d^5

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Electron configuration of Cu

[Ar] 4s^13d^10

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Effective nuclear charge

“charge” felt by other electrons. Ability to pull in electrons

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How to calculate Zeff

Z(# of protons) - S(screening electrons closer to the nucleus)

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Periodic table ->(Zeff)

Zeff increases as #of protons increases, but not the # of screening electrons

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Periodic Table goes down(Zeff)

Zeff is constant

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Periodic trends in atomic size

n and Zeff
As n increases, radius increases (PT goes down)
As Zeff increases, radius decreases (PT ->)

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Size of transition elements

their sizes do not change much

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Size of atoms vs cations

Cations(+) < atoms

Because Zeff increases

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Size of atoms vs anions

Anions(-) > atoms

Because Zeff decreases

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Trends in first ionization energies

closer is harder to ionize and vice versa. As Zeff increases leads to harder removal. As n increases, becomes easier to remove.

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2nd and higher ionization energie

I1 < I2 < I3 < …

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Ionic bonding

A cation and an Anion

electrons are given

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Covalent bond

Generally non-metal + non-metal | electron sharing

Metallic bonding

electrons are mobile in what is called an "electron sea" | conductors

Why do covalent bonds share electrons

To get lower in energy

Lattice energy

the energy it takes to separate one mol of ionic into infinite separation

Coulomb's law equation

E = (kQ1Q2)/R

Electronegativity

The force that an atom attracts electrons in a covalent bond.

As # of bonds increase

strength of the bond does too

Formal charge

Charge on an atom in a molecule by assigning one e- to each end of a single bond. Can be calculated by using original # of electrons - current #

Sum of the formal charges

charge on species. represents the charge of the molecule. Wants each formal charge to be as small as possible

Electron poor

When an atom is lacking an electron to achieve octet rule, but maintains 0 formal charge

electron rich

an atom has too many valence electrons, but still maintains 0 formal charge

odd electron

an atom has an unpaired valence electron, but still maintains 0 formal charge

Bond energy

it takes energy to break a bond and releases energy when a bond is formed. The average value can be found on the table at the back of the data sheet

molecular geometry

shapes of molecules and bond angles. It takes at least 3 atoms to make an angle

steric number (sn)

number of electron groups around an atom. Lone pairs + atoms attached to central atom

example of electron poor molecule

BF3

example of odd electron

example of electron rich

PF5

Angle and name of steric number 2

180 and linear

angle and name of sn 3

120 degrees and trigonal pyramid

name and angle of sn 4

109 degrees and tetrahedral

name and angle of sn 5

6 90 degrees, 2 120 degrees and 1 180 degrees, trigonal bypyramid

angle and name of sn 6

90 and 180 degrees, octahedral/square pyramidal

dipole moment and polarity

if a force of attraction from a polar bond creates a net vector that is not 0, the molecule is polar

Anion

negatively charged ion

cation

positively charge ion

Orbital hybridization

When orbitals intersect during bonding

Types of hybridization

``` SN 2: sp, p+p SN 3: sp2, p Sn 4: sp3 SN5: sp3d SN6: sp3d2 ```

Single bond type of hybridization

Sigma

Double bond type of hybridization

Triple bond type of hybridization

sigma and pi

Intermolecular forces

forces between molecules

Dipole dipole interaction

Polar molecules, the more polar it is, the higher the dipole-dipole force

London dispersion force

Depends on the number of electrons | the higher the molar mass, the higher the force

Hydrogen bonding

If H is connected to a small electro negative element(F,O,N), it is attracted to a lone pair on another electro negative element

Strength of each inter molecular forces

Hydrogen bonding > dipole - dipole > london dispersion force

boiling point related to inter molecular forces

The higher the inter molecular forces, the higher the boiling point

Isomer

Two molecules with the same molecular formula but different structures(Ie Ch3OCH3 and CH3CH2OH)

vapor pressure in relation to molar mass

the smaller the molar mass the higher the vapor pressure

vapor pressure compared to intermolecular forces

higher the forces, lower the vapor pressure

Vapor pressure in relation to temperature

vp increases as t increases