Unit 2 - Bonding and phases

Ionic bonds are usually formed between

Metals and non-metals

Electrons in Ionic Compounds are…

Localized around the anion as the cation gives up its electrons

Properties of Ionic substances are…

• Poor conductors of electricity

  • Due to lattice structures and localized electrons

  • Ionic Liquids can conduct electricity as the ionic molecules themselves move

• High melting points

  • Higher coulombic attraction = Higher melting point (vice versa)

• High Boiling points

Metallic bonds are formed between…

the same metallic element

Properties of metallic bonds…

• good conductors of electricity

  • the positive core (nucleus and core electrons) are stationary while the valence electrons move around

Alloy

combination of 2 different metals

Interstitial alloy

formed between metals of 2 different radii (smaller atom occupies the crevices “interstices” formed by the bigger element

Substitutional Alloy

Formed between atoms of similar radii (some atoms are substituted for atoms of the other)

Covalent bonds

the sharing of electrons by atoms

sigma (σ) bond

the first bond formed in a covalent bond

pi (π) bond

the second and third bonds formed in a covalent bond ( one pi bond, double pi bond)

Covalent bonds form at minimum potential energy because…

• Too close = the repulsive forces push away the atoms and no bond can form

• Too far = The nucleus of one atom cannot attract the electron of another

Network Covalent bonds

Elements that form lattice like structures with localized electrons

• Most common network covalent are either C or Si as they have 4 valence electrons

Doping

The process of increasing conductivity by adding an impurity

P-doping

some of the atoms in the structure is replaced with an atom with fewer valence electrons which leaves a hole in the bonding (positive) so it draws outside electrons that create a chain reaction to increase the overall conductivity

N-doping

some of the atoms in the structure is replaced with an atom with more valence electrons which have nowhere to bond so they roam around increasing the conductivity

The central element in a Lewis dot structure is…

• the least electronegative

• have atleast 8 to a max of 12 valence electrons

Resonance forms/ structures

used to show that electrons are delocalized and can be shared with multiple atoms in an molecule

In a resonance form all the bonds…

have the same length and strength (somewhere between a double bond and a single bond)

Bond order

The way to determine the strength and length of bonds

• single bonds have a bond order of 1

• double bonds have a bond order of 2

• triple bonds have a bond order of 3

You can calculate the bond order for resonance structures by…

• Add the total bonds

• Divide by the number of possible resonance forms

• Ex: \frac{\left(1+2+1\right)}{3}=1.33

Only these 3 elements do not need 8 valence electrons to be stable…

• H (Hydrogen) only needs 2

• He (Helium) only needs 2

• B (Boron) only needs 6

Formal Charge

as Lewis dot structures can be drawn in multiple ways this is the way to find the correct way to draw them

Rules of Formal Charge

• Neutral molecule has a formal charge of 0

• Polyatomic Ion has a formal charge = to the overall charge of the ion

How to calculate formal charge…

• subtract valence electrons from each atom from the assigned electrons

  • For assigned electrons

    • Lone pairs count as 2

    • Bonds count as 1

VESPR Theory (valence electron shell pair repulsion)

electron pairs repel each other so they position themselves to minimize repulsion

Electron groups

lone pairs, bonds, single upaired electron, etc

this has 2 electron groups

linear

this has 3 electron groups

trigonal-planar

this has 4 electron groups

tetrahedral

this has 5 electron groups

trigonal-bipyramidal

this has 6 electron groups

octahedral

terminal electrons

the outermost electrons in a molecule

• 0 lone pairs

• 180 degree angle

• Linear

• 0 lone pairs

• 120 degree angle

• Trigonal Planar

• 1 lone pair

• 120 degree angle

• Bent trigonal planar

• 0 lone pairs

• 109.5 degree angle

• tetrahedral

• 1 lone pair

• slightly less than 109.5 degrees

• Trigonal Pyramidal

• 2 lone pairs

• slightly less than 109.5

• Tetrahedral Bent

• 0 lone pairs

• Trigonal Bi-pyramidal

• 1 lone pair

• Trigonal Bi-pyramidal seesaw

• 2 lone pairs

• Trigonal bi-pyramidal t-shaped

• 3 lone pairs

• Trigonal bi-pyramidal linear

• 0 lone pairs

• Octahedral

• 1 lone pair

• square- pyramidal

• 2 lone pairs

• square-planar

Polarity

In a covalent bond electrons aren’t shared equally as higher electronegative elements have the electrons crowding around them

Polar covalent bonds

In a covalent bond between 2 different electronegative elements and one will pull more atoms

Non-polar covalent bonds

A covalent bond between 2 of the same elements so electrons are shared equally

Dipoles

in polar covalent bonds the electrons crowding in one region creates a negative pole and a positive pole

Rule of thumb for molecular polarity

a molecule is non polar if it’s central atom has no lone pairs which means the molecule is symmetrical

Exceptions to the rule of thumb for molecular polarity

• Hydrogen will always be a positive dipole as its electronegativity is low

• molecules in a square planar shape are usually nonpolar as the terminal atoms are in the same plane despite central atom having lone pairs

Intermolecular Forces (IMFs)

forces between covalently bonded molecules that have to be broken fot phase change

Dipole-Dipole Forces

The positive dipole is attracted to negative dipoles

• substances with only dipole dipole forces are usually liquid or gases at stp

Hydrogen bonds

A type of dipole-dipole force, when the positive dipole of a hydrogen atom in one molecule is attracted to a very electronegative atom of another molecule

• strong as Hydrogen has no electrons to shield due to the bonding

London Dispersion Forces

random movements of electrons create temporary dipoles and occurs in all molecules

• can be stronger than H bonds in rare occasions

  • more electrons = more LDF

  • more molar mass = more electrons

  • more molar mass = more LDF

IMFs ranked by strength

• Hydrogen bonds

• Dipole-dipole forces

• LDF

Vapor pressure

when molecules in a liquid which are always in random motion hit the surface of the liquid with enough KE to turn into a gas

• temp ∝ vapor pressure

• strong IMF = weaker vapor pressure

Solutes

the thing to be dissolved

Solvents

the thing that its to be dissolved in

polar solutes dissolve best in…

polar solvents

non-polar solutes dissolve best in…

non-polar solvents

Electrolytes

the ions in a solvent that dissolved an ionic solute (conductivity increases)

Chromatography

the process of passing a mixture through a medium to separate it

paper chromatography

a paper is suspended above a (polar or nonpolar) solvent and mixture is blotted on paper as the solvent climbs the substances at mixtures climb to different rates depending on polarity, polar solutes climb the furthest for polar solvents

R_f

used to calculate the attraction between solute and solvent

higher R_f = more attraction

(distance traveled by solute/ distance traveled by solvent)

Column chromatography

a column is filled with a stationary substance then its injected with analyte which is at the stationary phase then its injected with eulent which as it exits the column the more attracted the analyte molecules are the faster they exit the column as well

Analyte

the solution to be seperated

Eulent

another solution

Distillation

Takes advantage of substances having different boiling points

• boils mixture at x degrees C and only the substance with x degrees C evaporates and is collected

Kinetic Molecular Theory

KE=\frac12mv^2

Assumptions of Kinetic molecular theory

• temp ∝ Avg. KE

• there is no attractions between the molecules in the gas

• if the sample is a mixture of gasses they will all have the same Avg. KE

• Gas molecules are in constant elastic motion

• volume of the ideal gas is insignificant compared to the volume of the container

Maxwell-Boltzman diagrams

used to model the velocities of different gases at different temps, or same gas at dif temps

Effusion

the rate of which a gas will escape through microscopic holes in a container

• higher temp = higher rate of effusion

• same temp, lower molecular mass = higher rate of effusion

Ideal Gas Law

PV=nRT

Boyles Law

as pressure increases, volume decreases

Charles Law

as temp increases, volume increases

Daltons Law

the total pressure of a gas is the sum of all the partial pressures, the partial pressure of a gas ∝ to the percent of moles of that gas in the mixture

Gases deviate from ideal state when…

the temp or pressure gets too high or too low

When gases deviate from ideal state

• molecules stick

• the volume of the gas is significant

Density

D=m/v or molar mass = DRT/P